Just use the websites above for study ^^^. I Haven't had a chance to add anything here yet...

Matter • Matter is the substance of which physical objects are composed. • It can be solid, liquid or gas. Solids

Unit D: Structures and Forces

Overview: Structures can be found in both natural and human-constructed environments, serving a variety of purposes and taking a wide range of forms. In learning about structures, students investigate the properties of materials used, and test them under different loads and forces. They examine different ways that structural components are configured, analyze forces involved, and investigate resulting effects on structural strength and stability.

As part of their study, students also examine construction methods used in the past and the present, and learn how science and technology link together in developing safe and efficient designs that meet human needs.

Focusing Questions: How do structures stand up under load?

What forces act on structures, and what materials and design characteristics contribute to structural strength and stability?

Key Concepts The following concepts are developed in this unit and may also be addressed in other units at other grade levels.

The intended level and scope of treatment is defined by the outcomes below. − structural forms − material strength and stiffness − joints − forces on and within structures (loads and stresses) − direction of forces − deformation − structural stability − modes of failure − performance requirements − margin of safety Outcomes for Science, Technology and Society (STS) and Knowledge

Students will:

1. Describe and interpret different types of structures encountered in everyday objects, buildings, plants and animals; and identify materials from which they are made

• recognize and classify structural forms and materials used in construction (e.g., identify examples of frame structures, such as goal posts and girder bridges, examples of shell structures, such as canoes and car roofs, and examples of frame-and-shell structures, such as houses and apartment buildings)

• interpret examples of variation in the design of structures that share a common function, and evaluate the effectiveness of the designs (e.g., compare and evaluate different forms of roofed structures, or different designs for communication towers)

• describe and compare example structures developed by different cultures and at different times; and interpret differences in functions, materials and aesthetics (e.g., describe traditional designs of indigenous people and peoples of other cultures; compare classical and current designs; investigate the role of symmetry in design)

• describe and interpret natural structures, including the structure of living things and structures created by animals (e.g., skeletons, exoskeletons, trees, birds’ nests)

• identify points of failure and modes of failure in natural and built structures (e.g., potential failure of a tree under snow load, potential failure of an overloaded bridge)

2. Investigate and analyze forces within structures, and forces applied to them

• recognize and use units of force and mass, and identify and measure forces and loads

• identify examples of frictional forces and their use in structures (e.g., friction of a nail driven into wood, friction of pilings or footings in soil, friction of stone laid on stone)

• identify tension, compression, shearing and bending forces within a structure; and describe how these forces can cause the structure to fail (e.g., identify tensile forces that cause lengthening and possible snapping of a member; identify bending forces that could lead to breakage)

• analyze a design, and identify properties of materials that are important to individual parts of the structure (e.g., recognize that cables can be used as a component of structures where only tensile forces are involved; recognize that beams are subject to tension on one side and compression on the other; recognize that flexibility is important in some structures)

• infer how the stability of a model structure will be affected by changes in the distribution of mass within the structure and by changes in the design of its foundation (e.g., infer how the stability of a structure will be affected by increasing the width of its foundation)

3. Investigate and analyze the properties of materials used in structures

• devise and use methods of testing the strength and flexibility of materials used in a structure (e.g., measure deformation under load)

• identify points in a structure where flexible or fixed joints are required, and evaluate the appropriateness of different types of joints for the particular application (e.g., fixed jointing by welding, gluing or nailing; hinged jointing by use of pins or flexible materials)

• compare structural properties of different materials, including natural materials and synthetics

• investigate and describe the role of different materials found in plant and animal structures (e.g., recognize the role of bone, cartilage and ligaments in vertebrate animals, and the role of different layers of materials in plants)

4. Demonstrate and describe processes used in developing, evaluating and improving structures that will meet human needs with a margin of safety

• demonstrate and describe methods to increase the strength of materials through changes in design (e.g., corrugation of surfaces, lamination of adjacent members, changing the shape of components, changing the method of fastening)

• identify environmental factors that may affect the stability and safety of a structure, and describe how these factors are taken into account (e.g., recognize that snow load, wind load and soil characteristics need to be taken into account in building designs; describe example design adaptations used in earthquake-prone regions)

• analyze and evaluate a technological design or process on the basis of identified criteria, such as costs, benefits, safety and potential impact on the environment

8 Unit A: Mix and Flow of Matter (Science and Technology Emphasis) Overview:

The materials that we use—including natural and manufactured ones—often take the form of fluids.

Students learn that such diverse substances as air, natural gas, water and oil are fluids. In further investigations, they discover that many common household materials are aqueous solutions or suspensions in which the main component is water.

Students learn that the properties of individual fluids are important to their use, including such properties as density, buoyancy, viscosity and the fluid’s response to changes in temperature and pressure.

The particle model of matter is introduced to help students make a conceptual link between the nature of matter and the specific behaviour of fluids. Focusing Questions: What are fluids? What are they made of and how do we use them?

What properties of fluids are important to their use? Key Concepts The following concepts are developed in this unit and may also be addressed in other units at other grade levels. The intended level and scope of treatment is defined by the outcomes below.

− WHMIS symbols and nomenclature

− pure substances, mixtures and solutions

− solute and solvent

− concentration

− solubility and saturation points

− particle model of matter

− properties of fluids

− viscosity and flow rate

− mass, volume, density

− pressure

− buoyancy Outcomes for Science,

Technology and Society (STS) and Knowledge Students will:

1. Investigate and describe fluids used in technological devices and everyday materials

• investigate and identify examples of fluids in household materials, technological devices, living things and natural environments

• explain the Workplace Hazardous Materials Information System (WHMIS) symbols for labelling substances; and describe the safety precautions to follow when handling, storing and disposing of substances at home and in the laboratory

• describe examples in which materials are prepared as fluids in order to facilitate transport, processing or use (e.g., converting mineral ores to liquids or slurries to facilitate transport, use of paint solvents to facilitate mixing and application of pigments, use of soapy water to carry away unwanted particles of material)

• identify properties of fluids that are important in their selection and use (e.g., lubricant properties of oils, compressibility of gases used in tires)

2. Investigate and describe the composition of fluids, and interpret the behaviour of materials in solution

• distinguish among pure substances, mixtures and solutions, using common examples (e.g., identify examples found in households)

• investigate the solubility of different materials, and describe their concentration (e.g., describe concentration in grams of solute per 100 mL of solution)

• investigate and identify factors that affect solubility and the rate of dissolving a solute in a solvent (e.g., identify the effect of temperature on solubility; identify the effect of particle size and agitation on rate of dissolving)

• relate the properties of mixtures and solutions to the particle model of matter (e.g., recognize that the attraction between particles of solute and particles of solvent helps keep materials in solution)

3. Investigate and compare the properties of gases and liquids; and relate variations in their viscosity, density, buoyancy and compressibility to the particle model of matter

• investigate and compare fluids, based on their viscosity and flow rate, and describe the effects of temperature change on liquid flow

• observe the mass and volume of a liquid, and calculate its density using the formula d = m/v [Note: This outcome does not require students to perform formula manipulations or solve for unknown terms other than the density.]

• compare densities of materials; and explain differences in the density of solids, liquids and gases, using the particle model of matter

• describe methods of altering the density of a fluid, and identify and interpret related practical applications (e.g., describe changes in buoyancy resulting from increasing the concentration of salt in water)

• describe pressure as a force per unit area by using the formula p = F/A, and describe applications of pressure in fluids and everyday situations (e.g., describe pressure exerted by water in hoses, air in tires, carbon dioxide in fire extinguishers; explain the effects of flat heels and stiletto heels, using the concept of pressure)

• investigate and compare the compressibility of liquids and gases

4. Identify, interpret and apply technologies based on properties of fluids

• describe technologies based on the solubility of materials (e.g., mining salt or potash by dissolving)

• describe and interpret technologies based on flow rate and viscosity (e.g., heavy oil extraction from tar sands, development of motor oils for different seasons, ketchup/mustard squeeze bottles)

• describe and interpret technologies for moving fluids from one place to another (e.g., intravenous lines, pumps and valves, oil and gas pipelines)

• construct a device that uses the transfer of fluids to apply a force or to control motion (e.g., construct a model hydraulic lift; construct a submersible that can be made to sink or float by transfer of a fluid; construct a model of a pump) Skill Outcomes (focus on problem solving) Initiating and Planning Students will: Ask questions about the relationships between and among observable variables, and plan investigations to address those questions

• define practical problems (e.g., How can we remove a salt coating from a bicycle or vehicle?)

• identify questions to investigate, arising from practical problems and issues (e.g., identify questions, such as: “What factors affect the speed with which a material dissolves?”)

• phrase questions in a testable form, and clearly define practical problems (e.g., rephrase a question, such as: “Is salt very soluble?” to become “What is the most salt that can be dissolved in one litre of water at 23ºC?”)

• design an experiment, and identify the major variables (e.g., design or apply a procedure for measuring the solubility of different materials)

(grade 9) Unit B: Matter and Chemical Change

(grade 9) Unit B: Matter and Chemical Change (Nature of Science Emphasis) Overview: Different materials have different properties. The ability to distinguish between different substances and make sense of their properties, interactions and changes requires the development of ideas about chemical substance. In this unit, students are introduced to the formal study of chemical substance through laboratory investigations and introductory studies of chemical theory. In the laboratory, students observe and compare chemical substances and, with guidance on safety, investigate the properties of materials and the ways they interact. In conjunction with these studies, students are introduced to ideas about elements and compounds, and corresponding structural ideas about atoms and molecules. Theoretical ideas are introduced as means for explaining, interpreting and extending their laboratory findings; these ideas include a general introduction to the periodic table, chemical nomenclature and simplified ways of representing chemical reactions. This unit builds on: • Grade 8 Science, Unit A: Mix and Flow of Matter This unit provides a background for: • Science 10, Unit A: Energy and Matter in Chemical Change Focusing Questions: What are the properties of materials, and what happens to them during chemical change? What evidence do we have of chemical change; and what ideas, theories or models help us explain that evidence? Key Concepts The following concepts are developed in this unit and may also be addressed in other units at other grade levels. The intended level and scope of treatment is defined by the outcomes below. • Workplace Hazardous Materials Information System (WHMIS) and safety • substances and properties • endothermic and exothermic reactions • reactants and products • conservation of mass • factors affecting reaction rates • periodic table • elements, compounds and atomic theory • chemical nomenclature (introductory treatment) Outcomes for Science, Technology and Society (STS) and Knowledge Students will: 1. Investigate materials, and describe them in terms of their physical and chemical properties • investigate and describe properties of materials (e.g., investigate and describe the melting point, solubility and conductivity of materials observed) • describe and apply different ways of classifying materials based on their composition and properties, including: − distinguishing between pure substances, solutions and mechanical mixtures − distinguishing between metals and nonmetals [Note: Metalloids may also be introduced at this level but are not required.] − identifying and applying other methods of classification • identify conditions under which properties of a material are changed, and critically evaluate if a new substance has been produced Unit B: Matter and Chemical Change Grade 9 Science /57 ©Alberta Education, Alberta, Canada 2. Describe and interpret patterns in chemical reactions • identify and evaluate dangers of caustic materials and potentially explosive reactions • observe and describe evidence of chemical change in reactions between familiar materials, by: − describing combustion, corrosion and other reactions involving oxygen − observing and inferring evidence of chemical reactions between familiar household materials • distinguish between materials that react readily and those that do not (e.g., compare reactions of different metals to a dilute corrosive solution) • observe and describe patterns of chemical change, by: − observing heat generated or absorbed in chemical reactions, and identifying examples of exothermic and endothermic reactions − identifying conditions that affect rates of reactions (e.g., investigate and describe how factors such as heat, concentration, surface area and electrical energy can affect a chemical reaction) − identifying evidence for conservation of mass in chemical reactions, and demonstrating and describing techniques by which that evidence is gathered. 3. Describe ideas used in interpreting the chemical nature of matter, both in the past and present, and identify example evidence that has contributed to the development of these ideas • demonstrate understanding of the origins of the periodic table, and relate patterns in the physical and chemical properties of elements to their positions in the periodic table—focusing on the first 18 elements • distinguish between observation and theory, and provide examples of how models and theoretical ideas are used in explaining observations (e.g., describe how observations of electrical properties of materials led to ideas about electrons and protons; describe how observed differences in the densities of materials are explained, in part, using ideas about the mass of individual atoms) • use the periodic table to identify the number of protons, electrons and other information about each atom; and describe, in general terms, the relationship between the structure of atoms in each group and the properties of elements in that group (e.g., use the periodic table to determine that sodium has 11 electrons and protons and, on average, about 12 neutrons; infer that different rows (periods) on the table reflect differences in atomic structure; interpret information on ion charges provided in some periodic tables) [Note: Knowledge of specific orbital structures for elements and groups of elements is not required at this grade level.] • distinguish between ionic and molecular compounds, and describe the properties of some common examples of each 4. Apply simplified chemical nomenclature in describing elements, compounds and chemical reactions • read and interpret chemical formulas for compounds of two elements, and give the IUPAC (International Union of Pure and Applied Chemistry) name and common name of these compounds (e.g., give, verbally and in writing, the name for NaCl(s) (sodium chloride), CO2(g) (carbon dioxide), MgO(s) (magnesium oxide), NH3(g) (nitrogen trihydride or ammonia), CH4(g) (carbon tetrahydride or methane), FeCl2(s) (iron(II) chloride), FeCl3(s) (iron(III) chloride) 58/ Grade 9 Science Unit B: Matter and Chemical Change ©Alberta Education, Alberta, Canada • identify/describe chemicals commonly found in the home, and write the chemical symbols (e.g., table salt [NaCl(s)], water [H2O(l)], sodium hydroxide [NaOH(aq)] used in household cleaning supplies) • identify examples of combining ratios/number of atoms per molecule found in some common materials, and use information on ion charges to predict combining ratios in ionic compounds of two elements (e.g., identify the number of atoms per molecule signified by the chemical formulas for CO(g) and CO2(g); predict combining ratios of iron and oxygen based on information on ion charges of iron and oxygen) • assemble or draw simple models of molecular and ionic compounds (e.g., construct models of some carbon compounds using toothpicks, peas and cubes of potato) [Note: Diagrams and models should show the relative positions of atoms. Diagrams of orbital structures are not required at this grade level.] • describe familiar chemical reactions, and represent these reactions by using word equations and chemical formulas and by constructing models of reactants and products (e.g., describe combustion reactions, such as: carbon + oxygen → carbon dioxide [C(s) + O2(g) → CO2(g)]; describe corrosion reactions, such as: iron + oxygen → iron(II) oxide [Fe(s) + O2(g) → FeO(s)]; describe replacement reactions, such as the following: zinc + copper(II) sulfate → zinc sulfate + copper [Zn(s) + CuSO4(aq) → ZnSO4(aq) + Cu(s)]) [Note 1: This outcome does not require students to explain the formation of polyatomic ions. Some chemicals with polyatomic ions may nevertheless be introduced; e.g., a brief introduction to CuSO4(s), ZnSO4(s) and H2SO4(aq) can help prepare students for further study of these materials in units C and D.] [Note 2: At this grade level, students are not required to balance reactants and products in chemical equations. Teachers may want to inform students about opportunities for further study of chemistry in Science 10 and in Science 14–24.] Skill Outcomes (focus on scientific inquiry) Initiating and Planning Students will: Ask questions about the relationships between and among observable variables, and plan investigations to address those questions • identify questions to investigate (e.g., ask questions about the reactivity of particular materials or about conditions that affect the rate of reaction, after observing that materials react at different rates) • define and delimit questions and problems to facilitate investigation (e.g., reframe a general question, such as: “What affects the speed of reactions?” to become one or more specific questions, such as: “How will temperature affect the rate of reaction between materials x and y?” or “How will moisture affect the rate of reaction between x and y?”) • state a prediction and a hypothesis based on background information or an observed pattern of events • select appropriate methods and tools for collecting data and information and for solving problems (e.g., plan and conduct a search for information about chemical elements, using appropriate print and electronic sources) Unit B: Matter and Chemical Change Grade 9 Science /59 ©Alberta Education, Alberta, Canada Performing and Recording Students will: Conduct investigations into the relationships between and among observations, and gather and record qualitative and quantitative data • carry out procedures, controlling the major variables (e.g., investigate the effect of particle size on a chemical reaction, taking care to identify and control other potentially relevant variables) • observe and record data, and prepare simple drawings (e.g., represent a molecule studied through a drawing) • demonstrate knowledge of WHMIS standards, by using proper techniques for handling and disposing of laboratory materials • research information relevant to a given question (e.g., research properties of materials) Analyzing and Interpreting Students will: Analyze qualitative and quantitative data, and develop and assess possible explanations • compile and display data, by hand or computer, in a variety of formats, including diagrams, flow charts, tables, bar graphs, line graphs and scatterplots (e.g., present data on different chemical substances in a form that facilitates interpretation) • calculate theoretical values of a variable (e.g., predict the total mass of the products of a chemical reaction, based on the mass of the reactants used) [Note: In this example, students can apply the law of conservation of mass.] • identify and suggest explanations for discrepancies in data • state a conclusion, based on experimental data, and explain how evidence gathered supports or refutes an initial idea • identify new questions and problems that arise from what was learned (e.g., identify new questions, such as: “Why do different compounds containing the same elements behave differently?” or “How do atoms stick together in a molecule?”) Communication and Teamwork Students will: Work collaboratively on problems; and use appropriate language and formats to communicate ideas, procedures and results • receive, understand and act on the ideas of others (e.g., follow given safety procedures) • evaluate individual and group processes used in planning and carrying out investigative tasks (e.g., evaluate the relative success and scientific merits of different approaches to drawing and making models of molecules)

MATTER & chemical change (grade 9)

American chemical Society website: https://www.youtube.com/user/AmerChemSoc

PH indicator simulation: https://phet.colorado.edu/sims/html/ph-scale-basics/latest/ph-scale-basics_en.html

MATTER

Notes

MATTER :

Anything that has mass and occupies space exists as a solid, liquid, or gas.

How do you change from one state to another? The state of matter depends on temperature (heat energy)

DIFFERENT TERMS TO DESCRIBE CHANGES OF STATE:

• Melting

• Evaporation

• Condensation

• Freezing

• Sublimation

• Deposition

TO UNDERSTAND HOW SUBSTANCES DIFFER, YOU NEED TO OBSERVE THEIR PROPERTIES

All matter has two types of properties:

Physical properties:

• color

• shine (luster)

• melting temp, freezing temp, boiling temp

• density

• hardness

• solubility

• conductivity

• malleability

• ductility

Chemical properties

• Determines the way a material behaves in a chemical reactions

• Ex. reaction to oxygen, acids, heat, water, ability to burn.

Checkpoint 1

• When a substance undergoes a physical change, its state may be altered, but its chemical composition is the same ex. Ice cream melting!

• A chemical change is when 2 or more substances react and form 1 or more new substances ex. H₂ + O₂ forms H₂O

CATEGORIZING MATTER

PURE SUBSTANCE:

Is made of only 1 kind of matter and has a unique set of properties

• A pure substance is further classified into an element or a compound

ELEMENT:

Matter than cannot be broken down into any simpler substance.

• The most purest matter because it contains only one kind of atom

• Represented by a symbol of a capital letter, or a capital letter with lower case letters.

COMPOUNDS:

A substance formed by 2 or more different elements chemically fixed in proportion.

MOLECULE:

Similar to a compound but elements are held together via a molecular bond

A MIXTURE IS 2 OR MORE SUBSTANCES PHYSICALLY (NOT CHEMICALLY) COMBINED

4 main types of mixtures:

MECHANICAL MIXTURE (HETEROGENEOUS MIXTURE):

Some or all particles can be seen and be separated (Ex. Salad)

SOLUTIONS (HOMOGENEOUS MIXTURE):

Can not visibly separate the particles , as one is dissolved in another (Ex. coffee)

A substance dissolved in water is called an aqueous solution

SUSPENSION:

Mixture where tiny particles of one substance are held within another. Substances can be separated by centrifuging or filtration. (Ex. Salad dressing)

COLLOID:

Cloudy mixture, where the particles are suspended and difficult to separate (Ex. milk)

TYNDALL EFFECT - LIGHT PASSES THROUGH A SOLUTION, HOWEVER LIGHT IS DISPERSED IN A COLLOID

OBSERVING CHANGES IN MATTER

PHYSICAL CHANGE:

When a substance undergoes a physical change, its state may be altered, but its chemical composition is the same

To identify a physical change:

• You can separate the end products (reactants) to form the products again. Ex. Salt water, sand & rocks

• You are able to re-freeze or melt the product again. Ex. Ice cream, plastic mould

ALUMINUM IODINE REACTION

CHEMICAL CHANGE:

When 2 or more substances react and form 1 or more new substances (The new substances formed are completely different from the original reactants)

Four actions that identify a chemical change:

• Change in color

• Change of odour (if present)…safety!!

• Change of composition (ex. precipitate formed, gas released)

• Release or absorption of energy in the form of light or heat

Checkpoint 2

CONTROLLING CHANGES OF MATTER FOR HUMAN CONSUMPTION

FREEZE DRYING OF FOOD:

• Food is frozen to ice (physical change)

• Frozen food placed in pressure chamber to sublime ice (solid to gas) (physical change). Water gas is removed.

• To eat food, stir in hot water

EVOLVING THEORIES OF MATTER

HUMANS BEGIN MANIPULATING MATTER IN THE STONE AGE

8,000 B.C. Humans begin to settle in the Middle East. Metals have yet to be discovered, tools were mostly made of stone. Fire is used to cook, kilns are used to bake glass (silica) , ceramics (clay), and bricks (clay).

6,000 BC to 1,000 BC - Chemists (more like metallurgist) begin melting metals (smelting) of high value to humans. Gold was popular due to its properties. Tin and copper are melted together to create bronze. Early human civilizations commonly used seven metals:

• Gold (Au) 6000 B.C.

• Copper (Cu) 4200 B.C.

• Silver (Ag) 4000 B.C.

• Lead (Pb) 3500 B.C.

• Tin (Sn) 2750 B.C.

• Iron (Fe) 1500 B.C.

• Mercury (Hg) 750 B.C.

2500 BC - Greek philosophers realize that rock can be broken down into small pieces and into powder. The first theory arises that matter is made up of tiny particles.

1200 BC - Hittites (An empire in the Middle East) learns how to extract iron from rocks (ores). Mixing iron with carbon creates steel - a very hard material for armour, weapons, tools. Hittites also create coinage (Iron age begins)

Metals are not the only important chemicals, chemist mix liquids too! Many cultures used juices and oils in everyday life and rituals.

400 BC - Greek philosopher Democritus stated that each type of material was made up of a different type of ‘atomos’ which means indivisible. These different particles gave each material its own unique set of properties. By mixing different atomos, you could make new materials with their own unique properties.

350 BC - A popular greek philosopher Aristotle (among others) stated that everything was made of elements such as earth, air, fire, and water. Because Aristotle was well known and well respected, his description of matter was preferred over Democritus’s description until 1600 AD, even though he was dead wrong.

THE AGE OF ALCHEMY

YAY ALCHEMY!

350 BC - 1200 AD - The largest empire of the time rises and falls (Roman empire) plunging Europe into the dark ages (a lot of knowledge is lost). Meanwhile chemists of the Middle East practiced pseudo-science of mixing metals and naturally occurring chemicals ( ‘Alkimiya’- Arabic for ‘chemist’ )

1200 AD - Alchemy migrates to Europe. Alchemists try to change common metals into gold (using magic & simple experiments). Alchemists are not wholly interested in understanding the nature of matter. Overall, alchemist do not engage experiments using the scientific method (they rather find the philosopher’s stone), however they have contributed to useful lab tools from their practice. Ex. Beakers, filters, flasks, crucibles, retort.

1597 AD - the German alchemist Andreas Libau published ‘Alchemia’, a book describing the achievements of alchemists and how to prepare chemicals.

FROM ALCHEMY TO ATOMS

1400 - 1500s - Renaissance begins, scientists begin to rediscover and understand the nature of matter and change. Based their theories on observations and experimentation rather than guesses and assumptions (sounds like scientific method to me).

1600’s - Irish chemist Robert Boyle modernizes chemistry (hypothesizes and observes the nature of chemicals), this is separate from the ‘alchemy’ (mixing chemicals to find gold, magic oriented)

Boyle also experimented with gases, and came up with proof supporting 400 BC Democritus’ tiny particle theory. Boyle believed matter was composed of tiny particles with various shapes and sizes that grouped together to form other individual substances. He wanted to determine what each type of particle was.

1770’s - French scientist Antoine Lavoisier studied chemical interactions between hydrogen, oxygen, and carbon. He developed a system for naming chemicals so all scientists could use the same words and be on the same page. He was given the title of ‘Father of Modern Chemistry’. Neat fact: he had a an awesome, intelligent, and very young wife that helped him. However he had his head chopped off during the french revolution.

BILLIARD BALL MODEL

1808 - John Dalton (English scientist and teacher)- suggested matter was made up of elements that are pure substances that contain no other substances. He put forward the first modern theory of atomic structure:

• Each element is composed of a particle called an atom

• All atoms of a specific element have identical masses

• Different elements have different atoms of different masses

He developed the ‘Billiard Ball Model’ where atoms are solid spheres. Theory held true until the discovery of electrons in 1897.

RAISIN BUN MODEL / PLUM PUDDING

1897 - J.J. Thompson (British physicist), discovered negatively charged sub-atomic particles electrons. He proposed the plum pudding model.

The model ascertains that atoms are like positively charged ‘goop/pudding’ and negatively charged electrons are embedded in it like little plums (yes very weird analogy)

This plum pudding theory held true until electrons were found to be outside the nucleus.

Despite his analogy J.J. Thompson was correct about electron charges balancing out proton charges, therefore atoms have zero electrical charge.

ATOMS IN THE MODERN AGE

1904 - Hantara Nagaoka (Japanese physicist)- proposed an atomic model that resembled a mini solar system - planetary model. Center had a large positive charge and negative electrons circled the positive center like planets orbiting the sun.

Many did not agree with this model as they couldn’t explain it. The model existed until scientists realized a) nucleus is not massive b) electrons are not closely connected to the nucleus

RUTHERFORD MODEL

1907 - Ernest Rutherford (British scientist) working in Montreal won the Nobel Prize for work in radioactivity. Supported Nagaoka’s model but modified it saying electrons float around randomly.

His model suggest that atoms were empty spaces which positive particles could pass through with a positive central core (he called it a nucleus).

Calculated the nucleus to be 1/10000 size of an atom. Like comparing a small green pea in a football field!

Rutherford’s Discovery Was a huge contribution to atomic theory

BOHR MODEL

1913 - Niels Bohr (Danish Physicist) theorized that electrons move in a specific circular orbits (electron shells) and they jumped between shells by gaining or losing energy

Late 1920’s - James Chadwick (British Physicist) discovered the nucleus contains protons (+ charge) and neutrons (no charge). Protons and neutrons have same masses. Electrons are about 2000 times smaller than protons or neutrons. They weigh next to nothing!

Early 1930’s - Quantum Theory explains that electrons do not orbit, but exist in a ‘charged cloud’ surrounding the nucleus. These clouds encircle the nucleus in various shapes. Scientists realize it is close to impossible to know the exact location of each electron in an atom, therefore scientist use ‘clouds’ to roughly estimate where an electron might be.

QUANTUM THEORY - WAY BEYOND GRADE 9 SCIENCE

THE GRADE 9 VERSION OF AN ATOM!

Not to be patronizing but teaching quantum physics to grade nine students is a little, well, tough. Therefore we will be using Niels Bohr’s version of an atom to help you conceptualize what an atom is all about. A nucleus with electrons orbiting around in orbitals/shells.

ATOMIC STRUCTURE

All elements within the periodic table follow the same basic shape of an atom, they all contain protons, electrons, and neutrons. Different elements have a different number of protons, electrons, and neutrons.

Atom structure can be split into two different parts:

Nucleus - The center of the atom, it contains:

PROTONS:

Sub atomic particle that has a POSITIVE charge (+)

NEUTRONS:

Sub atomic particle that has NO charge (Ø)

Pro-tip: the mass of an atom is the number of protons added to the number of neutrons.

Orbitals/Shells - Electrons exist in orbits that circle around the nucleus. Orbitals can hold multiple electrons . How many orbitals does an electron have? Depends on the atom, this will make more sense later!

ELECTRONS:

Sub atomic particle that has a NEGATIVE charge (-)

Valence Orbital/Shell - The furthest orbital from the nucleus (electrons in this orbit are referred to as valence electrons)

PARTS OF AN ATOM

ATOMS! WHAT’S THEIR DEAL?

Atoms are all about stability, they really dislike instability and will do anything to become stable.

Atoms can achieve stability by completely filling or completely emptying their valence shells. This can be achieve in two ways:

• The atom gains electrons in its furthest orbital

• The atom loses electrons in its furthest orbital

BACK TO CHEMISTRY AND THE ELEMENTS

Before scientist knew about atoms they unknowingly referred to different atoms as elements (Aristotle started this ‘elements’ trend). Elements were considered naturally occurring basic components of all matter.

CHEMICAL ELEMENT:

A pure substance made up of single atoms (these atoms are identical in structure)

• Ex. A nugget of gold is essentially a collection of individual atoms hanging out together that all have the same proton number (47).

• Ex. A lump of coal (carbon) is essentially a collection of individual atoms hanging out together that have the same proton number (6).

EARLY CHEMISTS LOOKED FOR PATTERNS AND PROPERTIES OF DIFFERENT ELEMENTS

1808 - John Dalton abbreviates names of the known elements, this is the beginning of using symbols rather than names for elements.

1814 - Jöns Berzelius (Swedish chemist) uses the first letter (capitalized) of the element name as the symbol (Ex. Hydrogen = H). If there were more than two elements starting with the same alphabetical letter- use a small second letter behind the capital (Ex. Helium = He)

1864 - John Newlands (English chemist) noticed a pattern when elements were listed by increasing atomic mass.

1869 - Dmitri Mendeleev (Russian chemist & card player) organized the known elements according to patterns in the properties of the elements. He showed that properties of elements repeat periodically with increasing atomic masses.

He charted these known elements into a table (63 elements were known at the time). This table would be the beginning of the periodic table of elements. His table had ‘gaps’ for future elements to be discovered that had properties and atomic masses to fit those gaps. He did this without even knowing what an atom was made of!

Dmitri Mendeleev noticed that elements in the same column acted similarly...strange?!

2020 - Since Dmitri Mendeleev introduced the periodic table of elements scientist have found all 118 elements.

THE PERIODIC TABLE…PAST AND PRESENT

Mendeleev’s periodic table had 63 elements. Since then, many more elements have been discovered. Mendeleev’s table also had holes in it…to help place a spot when new elements were discovered. These holes were placed based upon the patterns noticed.

Today, a Periodic Table has 118 elements. Some of these elements are very unstable and discovered using specific laboratory conditions.

Checkpoint 3

THE PERIODIC TABLE

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All the elements of the periodic table are arranged in a logical sequence. This sequence is the addition of one proton to each successive element. This arranges all the elements in order of increasing atomic weight. This method also arranges the elements according to similar properties (groups).

HOW TODAY’S PERIODIC TABLE IS ORGANIZED

PERIODS:

• 7 horizontal rows

• Left side is more chemically reactive than the right side, which is less chemically reactive

GROUPS:

• 18 vertical columns

• All elements in group columns have similar appearance & properties

ATOMIC NUMBER:

Number above the symbol. Indicates # of protons or electrons in an electrically neutral atom. Atomic Numbers increase by one from left to right within the ‘periods’

ATOMIC MASS:

Number below the name. Total mass of all protons and neutrons in the nucleus. Average mass of element’s atoms

Atomic mass measured in amu’s - atomic mass units where one amu = 1/12 mass of a carbon-12 atom. Each has a mass of 1 amu.

GENERAL GROUPINGS WITHIN PERIODIC TABLE OF ELEMENTS

THE GENERAL GROUPINGS ARE:

METALS:

Shiny, malleable, ductile, conduct electricity

Found on the left side over to the right past middle (Lithium 3 to Polonium 84)

METALLOIDS:

Have both metallic and non-metallic properties

Found between metals and non-metals

NON-METALS:

Solid (dull & brittle) or gases, don’t conduct electricity (insulators)

SPECIFIC GROUPINGS WITHIN THE PERIODIC TABLE OF ELEMENTS

ALKALI METALS:

Group 1 column has the most chemically reactive with air or water. Reactivity increases as you go down the group (NOT INCLUDING H)

ALKALINE-EARTH METALS:

Group 2 elements also react with air or water, but not as vigorously as group one

TRANSITION METALS:

Group 3 to group 12 metals that are found in the middle of the periodic table. They all have similar properties.

METALLOIDS:

Group 13 to Group 17 elements that are found under the staircase! Elements found here have properties that can be like typical metals and nonmetals

HALOGENS:

Group 17 elements are the most reactive non-metals. F, Cl, Br all react readily with Group 1 elements to produce useful compounds…like salt!

NOBLE GASES:

Group 18 elements are very stable and non-reactive. These gases combine with other elements under specialized lab techniques

CALCULATING THE NUMBER OF PROTONS, ELECTRONS, & NEUTRONS

• Proton number is the exact same as the atomic number, and it never changes numbers!

• Electron number is the exact same as proton number (in a ‘neutral’ atom, electron number can and will change, we’ll figure that out later)

• Neutron number depends on the atomic mass!

CALCULATING THE AMOUNT OF NEUTRONS

• Find the atomic mass and atomic number of the element

• Round off the atomic mass to the nearest whole number

• Subtract the atomic number from the rounded atomic mass

Ex. Find the amount of neutrons for Beryllium (Be)

• Atomic mass = 9.012 , Atomic number = 4

• Atomic mass rounded (9.012 rounds to 9)

• Atomic mass rounded (9) - atomic number (4) = 5

There are 5 neutrons per atom of Be

• P.E.N for Be is P = 4, E = 4, N = 5

DRAWING A SODIUM ATOM THE NIELS BOHR WAY

A SODIUM ATOM MODEL DRAWN HOW NIELS BOHR ENVISIONED IT!

Recall how Niels Bohr envision his atoms:

• Electrons orbit the nucleus in different orbital ‘shells’. These electron shells along with the nucleus create the whole structure that we call an atom.

• Atoms love being stable! Atoms are striving to either empty or fill their valence shell!

HOW TO DRAW A ‘NEUTRAL’ ATOM:

1. Determine PEN.

2. Place P & N in nucleus.

3. Place electrons outside nucleus (2 e-, 8 e-, 8 e-, 18 e- )

NOTE* we are drawing atoms that have a zero net charge, aka it is neutral and has no charge! (We will discuss charges in the next unit)

ISOTOPES

All atoms are not created equally! Some atoms can have a slightly different nucleus. The only difference is the neutron amount.

ISOTOPES:

These are atoms of the same element but have different numbers of neutrons in the nucleus. Their mass numbers will be different but atomic number is the same.

For instance the carbon atom has three different isotopes.

• ¹²C also known as Carbon-12 (most popular type of carbon on Earth, 99% of the time we find carbon it is Carbon-12 )

• ¹³C also known as Carbon-13

• ¹⁴C also known as Carbon-14 (a rare type of carbon) used for carbon dating

ATOMIC CHARGES

A SODIUM ATOM WHICH CONTAINS ONE ELECTRON IN ITS VALENCE SHELL - IT’S UNSTABLE!

Remember that all atoms are striving to become stable! They become stable by having their valence shell full or empty of electrons.

Why do we call the elements neutral?

Thus far we have discussed atoms in the periodic table (aka elements) which are in their neutral form. The atoms contain the same amount of electrons as protons. They have not lost or gained any electrons. But, they are unstable!

Atoms want to be stable, to achieve this they will need to either gain or lose electrons in their valence shell, only then will atoms be considered stable.

When atoms lose or gain electrons it affects their net charge.

ATOMIC NET CHARGE:

The net charge is equal to the proton number (+) added to the electron number (-) of an atom. The overall charge of an atom can be either negative or positive.

Ex. A sodium atom that has lost an electrons is (Na+)

Ex. A chlorine atom that has gained an electron is (Cl-)

What is the net charge of Sulfur that has gained 2 electrons? What is the net charge of Scandium that has lost 3 electrons?

IONS

Thus far we have been dealing with atoms which have the same number of electrons as protons.

To become more stable, atoms will try to empty or fill their farthest orbital (aka valence shell) with electrons. When atoms lose or gain electrons we call them ions.

ION:

Atoms that acquire an electrical charge through the loss or gain of electrons

If the atom loses electrons then it will acquire a positive charge (+)

If the atom gains electrons then it will acquire a negative charge (-)

There are special names for positive and negative ions:

CATION:

Atoms that lose electrons transform into positive ions (ex. Fe 2+, iron atom that has lost 2 electrons). Cations are attracted to negatively charged ions.

ANION:

Atoms that gain electrons transform into negative ions (ex. Cl 1-, chlorine atom that has gained 1 electron). Anions are attracted to positively charged ions.

To indicate an atom is an ion we place the charged number in superscript:

• A calcium atom that lost 2 electrons would be Ca²⁺ (also know as a calcium cation)

• A scandium atom that lost 3 electrons would be Sc³⁺ (also know as a scandium cation)

• A sulfur atom that gained 2 electrons would be S²⁻ (also know as a sulfur anion)

• A nitrogen atom that has neither gained nor lost electrons would be N⁰ - atoms that have a zero net charge are neutral, we do not indicate neutral/zero charges so in this case nitrogen would be indicated as plain old N.

If an atom has no charge is it called an ion?

IONIC CHARGE PATTERNS:

• Alkali metals (group 1) have 1+ charge

• Alkaline Earth metals (group 2) have a 2+ charge

• Group 16 have mostly a 2- charge

• Halogens (group 17) have a 1- charge

• Noble gases (group 18) have no charge

SPECIAL KINDS IONS

Thus far we have been discussing single individual ions such as Be²⁺, the fancy scientific name for individual ions is monatomic ion.

MONATOMIC ION:

An individual atom that has a positive or negative charge (ex. Ba+ or I-).

There are a few exceptions though! When atoms buddy up and act as one single unit, such as (PO₄), we use the fancy scientific name of polyatomic ion.

POLYATOMIC ION:

A group of atoms that behave as a single unit with a positive or negative electrostatic charge, some examples include:

• (PO₄)³⁻

• (SO₃)²⁻

• (H₃O)⁺

There is a list of polyatomic ions on the back of your Periodic Table of Elements sheet! Don’t let the complexity scare you!

COMPOUNDS

Where do atoms lose or gain electrons from? The answer is other atoms! Atoms are constantly swapping electrons to become more stable.

When atoms ‘trade’ or ‘share’ electrons they essentially bond to each other, this ‘buddying up’ is how compounds are formed.

COMPOUND:

When two or more elements chemically combine (aka atoms share or transfer electrons) they form a compound.

• When atoms ‘trade’ electrons they form an ionic bond, we refer to the ‘buddy’ grouping as an ionic compound. (ex. NaCl)

• When atoms ‘share’ electrons they form a covalent bond, we refer to the ‘buddy’ grouping as a molecular compound. (ex. H₂O)

IONIC COMPOUNDS

Remember when atoms interact with other atoms they may exchange electrons to become stable, and ‘buddy up’ to form compounds. Ionic compounds are formed when electrons are traded.

HOW A LITHIUM ATOM BONDS WITH A FLUORINE ATOM TO FORM LITHIUM FLUORIDE.

IONIC COMPOUND:

When two or more atoms transfer electrons they bond together. This bond creates an ionic compound and forms as a result of an attraction between particles of opposite charges (just like a magnet).

Ionic bonds involve cations (metals) donating/losing electrons to anions (non-metals).

Metals love donating electrons rather than sharing them.

FORMING IONIC COMPOUNDS

METAL ATOMS BOND TO NON-METAL ATOMS!

Remember that ionic compounds are formed by one type of atom donating/losing an electron to another atom. The charge on elements considered to be metals are always positive (+) while the charge on non-metals are mostly negative (-). Thus metals love bonding with non-metals.

FORMATION OF SODIUM FLORIDE (NaF)

Some example of ionic compounds include:

• Na⁺ bonds to F⁻ to form NaF (added to toothpaste to reduce cavities)

• K⁺ and Br⁻ to form KBr (antisiezure drug)

• Na⁺ and Cl⁻ to form NaCl (table salt)

• Fe²⁺ and O²⁻ to form FeO (rust)

CHARACTERISTICS OF IONIC COMPOUNDS

Characteristics that all ionic compounds have in common are:

• Ionic compounds are commonly known as ‘salts’ (this is different from table salt)

• Ionic compounds form hard and brittle substances

• Ionic compounds formed are solid at room temperature

• Ionic compounds have high melting point and boiling point

• Ionic compounds are good electrical conductor in an aqueous solution (dissolved in water)

CHEMICAL FORMULA TERMINOLOGY

Through hundreds of years compounds have been studied and chemicals have been named according to the scientist. Until the 18th century, no standardized system existed for naming chemicals.

In the 1780’s Guyton de Morveau began using the chemical symbol for each element to write out compounds. He decided the metal element in the compound is named first!

In 1920 IUPAC (International Union of Pure & Applied Chemistry) created the terminology for naming compounds and is now responsible for naming all chemical compounds discovered.

HOW TO UNDERSTAND CHEMICAL FORMULAS

Three components of a chemical formula include:

1. Chemical symbols - identifies what element(s) are present in the compound

• Ex. NaCl: one atom of sodium and one atom of chlorine

2. Subscript Numbers - indicates the number of atoms of elements that must combine to form the compound. It is placed below and after the symbol.

• H₂O - 2 hydrogen, 1 oxygen

• C₆H₁₂O₆ - 6 carbon, 12 hydrogen, 6 oxygen

3. Indicating Physical State of Element or Compound. After the chemical formula, as subscripts put: (s) for a solid compound, (l) for a liquid compound, (g) for a gaseous compound, (aq) compound dissolved in water

• C₆H₁₂O₆ ₍s₎ or H₂O₍ʟ₎

Checkpoint 1

Mini Review !

WRITING IONIC COMPOUNDS

SALT!

NAMING IONIC COMPOUNDS

1. The chemical name of the metal (cation) comes first followed by the name of the non-metal (anion), the charges of each element are not written.

• Ex. CaO is calcium oxide

2. Name of the non-metal (anion) changes to an ‘ide’ suffix.

• Ex. Al2O3 - aluminum oxide

• Ex. TiS₂ - titanium sulfide

* An exception to the ‘ide’ rule is if the anion is polyatomic, then the name of the polyatomic ion is used

• Ex. CaSO₄ - calcium sulfate

3. If an element has multiple charges (ex. copper has a 1+ or 2+ charge), a Roman Numeral is added after the cation (Metal).

• Ex. vanadium (IV) oxide

• Ex. Iron (III) nitride

• Ex. copper (II) Sulfate is CuSO4

FORMING IONIC COMPOUNDS

ALWAYS have a copy of the Periodic Table

1. Print the metal element’s symbol with its ion charge, next to it, print the non-metal element’s symbol with its ion charge.

• Ex: Ca2+ Cl1-

2. Balance the ion charges, the positive ion charge must balance the negative ion charges

• Ex: Ca2+ Cl1- Cl1-

In the example, this means that there must be two chlorine atoms each with an ion charge of 1- to balance the 2+ ion charge of one calcium atom.

3. Write the formula by indicating how many atoms of each element are in it.

• Do not include the ion charge in the formula!

• Place the number of atoms of each element in a subscript after the element’s symbol.

• If there is only one atom, no number is used

  • Ex:CaCl2

The Criss Cross Method!

Naming Ionic Compounds Worksheet

Checkpoint 2

MOLECULAR COMPOUNDS

Thus far we have discussed how atoms can ‘trade’ electrons to become stable and form compounds. What if instead of trading they ‘shared’ electrons.

With molecular compounds electrons are shared. This means that electrons are neither lost nor gained by atoms. This also means atoms do not become ions! What elements share electrons? The non-metals are all about sharing!

MOLECULAR COMPOUND:

When one or more non-metallic atoms combine they form a molecular compound (ex. CO, CO₂, H₂O). Another name for molecular compound is molecule.

• Rather than donating electrons (like ionic compounds) molecular electrons share electrons, by sharing electrons atoms form a covalent bond.

• There are no cations and anions when we discuss molecular compounds!

CHARACTERISTICS OF MOLECULAR COMPOUNDS

Molecular compounds characteristics differ greatly from ionic compounds characteristics:

WATER IS A MOLECULAR COMPOUND

• Can be solids, liquids, or gases at room temp.

• Poor electrical conductors

• Low melting and boiling points

• All molecular elements are non metallic

DIATOMS

THE OXYGEN WE BREATHE IS DIATOMIC

Thus far we have discussed polyatomic ions (more than one atom), and monatomic (only one atom).

Remember non-metals love, love, love to share electrons. In fact many non-metals ‘buddy up’ with each other to form a somewhat stable pair. These molecules are considered to be diatoms.

DIATOMIC ELEMENTS:

Elements that form molecules consisting of two atoms bonded together.

All halogens are diatomic, as well as N, H, and O:

• Ex. Oxygen - O₂

• Ex. Hydrogen - H₂

• Ex. Chlorine - Cl₂

Molecules that are made of the same atom (ex. O₂) are homonuclear, while molecules made of different atoms are consider to be heteronuclear (ex. H₂O).

NAMING MOLECULAR COMPOUNDS

Similar to writing ionic compounds except:

• No ions are present!

• Ion charge not used in the formulas

Formula will show: What elements are present & proportion of atoms in a molecule

• Ex. O₂₍g₎ has 2 oxygen atoms

IUPAC recommends that molecular compounds should be named using the prefix system only.

In the prefix system, Greek or Roman prefixes are used to indicate the number of each kind of atom bonded to one another

MOLECULAR RULES:

• The first element in the formula should be named in full.

• The second element in the formula should be shortened and given an ‘ide’ suffix…just like ionic compounds!

• When there is more than one atom in the formula, a prefix is used which specifies the number of atoms.

• The prefix ‘mono’ may be omitted on the first element

FORMING MOLECULAR COMPOUNDS

Predicting how molecular compounds will form is a much more complex process than with predicting ionic compounds, take for instance sugar C₁₂H₂₂O₁₁. How non-metals combine to form molecular compounds can be tricky to understand and is out of scope for grade 9 science. It’s best to save this lovely topic of molecular compounds for high school chemistry ;)

CHEMICAL REACTIONS

KABLAMO! POTASSIUM NITRATE, CARBON, AND SULFUR COMBINE TO FORM GUN POWDER - AN INGREDIENT OF FIREWORKS

RECALL THE EVIDENCE OF A CHEMICAL REACTION:

• Color change

• Odour

• Gas or solid formation

• Release or absorption of energy

CHEMICAL REACTIONS :

When two or more ionic or molecular substances combine to form new substances

• Atoms are rearranged

• New bonds are formed

• New substances with new properties form

EXOTHERMIC REACTIONS:

Chemical reactions that release heat.

• Ex. Burning wood releases energy in the form of heat and light

ENDOTHERMIC REACTIONS:

Chemical reactions that absorb energy. reactions need a continuous supply of energy.

• Ex. A cold pack is triggered by a chemical reaction that absorbs heat from its surrounding.

WE CAN EXPRESS CHEMICAL REACTIONS IN 3 DIFFERENT WAYS…

• Naming the chemical reaction: combustion of coal

• Using a word equation: carbon + oxygen → carbon dioxide

• Using a formula equation: C + O₂ → CO₂

WRITING A FORMULA EQUATION

Chemical reactions are often written as a formula equation. There are a few technical terms to know about formula equations:

• Reactants: materials at the start of a reaction

• Products: new materials produced by the reaction

• An arrow indicates direction of reaction

  • A + B → C + D

  • (Reactants) → (Products)

• A plus (+) sign is used to separate different reactants or products

Lets use the chemical equation for photosynthesis to illuminate the terms of a chemical equation

carbon dioxide + water → glucose + oxygen

CO₂ + H₂O → C₆H₁₂O₆ + O₂

• The names of the reactants are written to the left of the arrow. The names of the products are written on the right of the arrow. Compounds are separated by a plus (+) sign.

• The arrow can be read as “yields”, “forms” or “reacts to produce”

THE THREE C’S OF OXYGEN

Because oxygen is readily available and a very good electron acceptor it tends to be used be a common element used in chemical reactions.

• Combustion (burning) – oxygen reacts with a fuel to produce carbon dioxide and water

  • Fuel (ex. wood) + oxygen → carbon dioxide + water

• Cellular respiration (burning of sugar) - occurs in aerobic organisms (consumers that needs oxygen)

  • Food (ex. glucose) + oxygen → carbon dioxide + water

• Corrosion (rusting) - a slow chemical change involving oxygen with some metals

• Metal (ex. iron) + oxygen → rusted iron

CONSERVATION OF MASS IN CHEMICAL REACTIONS

CHEMICAL REACTION IN AN OPEN SYSTEM - NO STOPPER ON THE FLASK!

OPEN SYSTEM:

An experiment in which the exchange of matter as well as energy interacts with the surroundings

CLOSED SYSTEM:

An experiment in which all reactants and all products of a chemical reaction are accounted for

LAW OF CONSERVATION:

Matter cannot be destroyed but can change form in a ‘closed’ controlled laboratory system.

Reactants → Products

100 g → 100 g

Mg + S → MgS

24.3g + 32.1 g → 56.4 g

Matter is not created nor destroyed during a chemical rxn. However, this being said, some matter such as gases may and usually escape and not be accountably measured in an ‘open system’ reaction as performed in most school labs

Conservation of Matter Worksheet

BALANCING EQUATIONS

All chemical reactions must be balanced—the number of atoms, moles, and ultimately the total mass must be conserved during a chemical process.

1. Count the number of atoms on both the reactants and product sides of the equation

2. Using coefficients, balance the number of atoms. Always leave diatoms such as hydrogen and oxygen last

3. Check your work by counting the number of atoms on each side of the equation

Balancing Equation Worksheet

SPEED OF CHEMICAL REACTIONS

Chemical reactions can occur at different rates, take for instance rusting of metal which takes days, versus a firework which reacts within a matter of seconds. Certain factors can affect the speed that a reactions occurs at.

Recall that chemical reactions are essentially atoms interacting with each other, anything that increases the amount of interaction between atoms will increase reactions rates.

FOUR FACTORS AFFECTING RATE OF REACTION:

1. The presence of catalysts

CATALYSTS:

Any substance that increases the rate of a reaction without itself being consumed.

• catalysts speed up reaction

• not used in the reaction itself

• enzymes in the body are a great examples of catalysts

2. Concentration of reactants

• The greater the concentration, the faster the reaction. Why? More atoms to react with each other

3. Temperature of the reactants

• Higher the heat, faster the reaction. Why? Causes the atoms to move faster creating more collisions with each other

4. Surface area of reactants

• Greater surface area (ex. Powder vs chunks) the more surface area will be available for the reaction. More surface area means that more area is available in relation to the volume for reaction.