Particles

Essential Ideas: 

Chemists develop and use models to understand the atomic world

Atoms combine to form molecules and compounds through various types of chemical bonds

Everything is made up of atoms and those atoms react to form molecules 

Learning Objectives:

Atomic structure (most of these learning objectives are revision of concepts learned in Year 8)

• Use the symbols of the elements correctly

• Explain that the Periodic Table orders elements horizontally by the number of protons in the atom’s nucleus and places those with similar chemical properties in columns.

  • Model how each atom has a structure consisting of a nucleus, which is made of protons and neutrons (nucleons), surrounded by electrons. Electrons are located in energy levels.

  • Deduce the number of valence electrons from an element's group (column) number

    • Describe how this leads to the similar properties of elements of the same group.

  • Describe the build-up of electrons in ‘shells’/’energy levels’ and how this structure is illustrated in the Periodic Table (periods/row number correspond to the number of occupied shells/energy levels)

• Understand that the mass of protons and neutrons are similar and determine the atomic mass. Electrons are orders of magnitude less massive.

  • Explain that nucleon number (protons + neutrons) is equivalent to atomic mass

  • Describe that atoms of the same element can have different numbers of neutrons and that these are called isotopes

  • Deduce the mass of the most common isotope based on the relative atomic mass given on the Periodic Table

    • Extension - Calculate the relative atomic mass based on the abundance of the regularly occurring isotopes (weighted average)

  • First introduction of moles

    • Describe how the scale of an atom is so small that a collective quantity, a mole, is used

    • Explain that a mole worth of atoms scales up the unit of atomic mass to grams. 

    • Explain that moles allow the comparison of a relative number of particles.

      • Compare masses of samples of elements, determine relative number of atoms (moles)  in each sample

Bonding

• Describe the difference between an element, mixture, and compound

• Describe the significance of the noble gas electronic structures and of valence electrons

  • Having the number of valence electrons of the nearest noble gas (group 18) results in a more stable structure

  • To achieve this stability, valence electrons can be lost, gained, or shared

Metallic Bonding

• Describe that metallic bonding involves atoms that are weakly attracted to their valence electrons.

• State that metallic bonding is the electrostatic attraction between metal cations and delocalized electrons.

• Describe how, in metal solids, atoms are arranged in a lattice.

• Describe how the weak attraction for their valence electrons results in the valence electrons moving free, to be delocalized, throughout the lattice structure.

• Explain how metallic bonding leads to the physical properties of metals

Ions and ionic compounds

• • Describe the formation of ionic bonds between metal and nonmetal elements

    • Describe ions as particles with more or less electrons than protons.

      • The number of protons don’t change, only the number of electrons

      • Most elements rarely exist in their atomic form. 

    • Describe metals as atoms that are weakly attracted to their valence electrons

      • There are few number of protons relative to the distance between the nucleus and valence shell/energy level

      • Losing valence electrons leads to the stable configuration of the preceding noble gas

      • Fewer electron than protons = positive ion called a cation

    • Describe nonmetals as atoms are strongly attracted to additional electrons

      • There are many protons relative to the distance between the nucleus and valence shell/energy level

      • Gaining valence electrons leads to the stable configuration of the subsequent/following noble gas

      • More elections than protons = negative ion called anion

    • Explain that there are no ionic elements, only ionic compounds

    • Define an ionic bond as the electrostatic attraction between oppositely charged ions

      • In ionic solids the ions are in a 3D lattice of alternating positive and negative ions

        • NOT MOLECULES!!!

    • Determine the formula of an ionic compound from the charges of the ions

    • State that ions of transition metal elements can have different charges

      • Use the roman numeral naming convention to determine charge and formula

        • i.e. iron(II) is Fe2+, iron(III) is Fe3+. Therefore the formula of iron(II)oxide is FeO and iron(III)oxide is Fe2O3 (good examples to start with as they avoid misconceptions with roman numerals indicating subscript).

  • Deduce the name or formula of ionic compounds containing the following polyatomic ions

    • Carbonate (CO32-), hydrogencarbonate (HCO3-), nitrate (NO3-), sulfate (SO42-), hydroxide (OH-), ammonium (NH4+)

  • Calculate the molar mass of various ionic compounds

    • Compare masses of samples of ionic compounds, determine relative number of formula units and atoms (in moles) in each sample

Covalent bonds, molecular compounds, and molecular elements

• • Define a covalent bond as the electrostatic attraction between two positive nuclei and a shared pair of electrons

  • Describe the formation of a covalent bond

    • Electron pairs can be shared when both atoms’ nuclei are strongly attracted to additional electrons

  • Describe with the support of a diagram the covalent bonds in a range of molecules including H2, Cl2, H2O, CH4, HCl, N2, NH3, C2H4, O2, CH3OH and CO2 as the sharing of pairs of electrons leading to the noble gas configuration. Include description of single, double, and triple bonds.

    • Deduce Lewis structures of a range of molecules

      • Use the octet tool (tool in that it works for explaining bonding in most, but not all, molecules. Examples of exceptions need not be discussed)

      • Use of dot cross diagrams should be avoided

        • Translate/transition Bohr to Lewis

  • Calculate the molar mass of various covalent compounds

    • Compare masses of samples of covalent compounds, determine relative number of molecules and atoms (in moles) in each sample

Structure and physical properties

• • Model how solids may be formed from molecules, or they may be extended structures with repeating subunits (e.g. ionic lattice, macromolecules)

    • In ionic solids the ions are in a 3D lattice of alternating positive and negative ions

    • In solids, molecules are held together by relatively weak attractions between the molecules called intermolecular forces

    • Macromolecules/Giant covalent structures can be introduced but won’t be assessed

  • Use differences in bonding to explain, in general, the physical properties of various elements (metals and nonmetals) and compounds (ionic and covalent). 

    • Particles should not be represented as only circles

      • Diagrams of ionic solid should show charge

      • Diagrams of molecular solid should imply that the molecules are made of multiple atom

    • All ionic compounds are solids at room temperature

    • Ionic solids are hard/brittle and do not conduct electricity (charged particles can’t move)

    • As a liquid (molten), the ions can move by one another therefore can conduct electricity

    • Ionic solutions conduct electricity because they have charged particles that can move

    • Molecules can be found as solids, liquids, or gasses at room temperature