Naming Review

• Ionic

  • Naming - cation goes first, anion second, change the ending to “ide”

  • Formulas - write cation and anion symbols and charge, cross over and reduce

• Covalent

  • Naming - Use the correct prefix for the number of atoms

  • Formula - Use the correct number of atoms according to the prefixes

• Acids

  • Naming - Hydro-”atom”-ic acid or “polyatomic atom”-ic acid

  • Formula - Hydrogen is in front, and then the other atoms

• Multivalent

  • Naming - reverse cross-over, use Roman numerals to state the charges

  • Formulas - apply indicated charge

• Hydrate - A compound that has associated with water

  • Use numerical prefix to indicate how many water molecules are present

  • Use • to show that water is associated with the molecule

• Polyatomic - A group of atoms with an atomic charge

  • Naming - cation first, anion second

  • Formulas - write the charge of both and cross over, but do not change the formula of the polyatomic ion

    • Oxyanion - A polyatomic ion with oxygen

      Use following endings for oxyanions:

      # of O atoms	Prefix	Suffix
      +1 O atom	per-	-ate
      Common #	-	-ate
      -1 O atom	-	-ite
      -2 O atom	hypo-	-ite

Lesson 1 - Average Atomic Mass

• The nucleus of an atom contains protons and neutrons

• Proton - Positively charged subatomic particle

  • Gives the atom its identity as an element

  • Atomic Number - The number of protons in an atom

  • They contribute to the atom’s mass

• Neutron - A subatomic particle that does not carry a charge, but contributes to an atom’s mass

  • To find the number of neutrons: atomic mass - atomic number

• Electrons - Negatively charged subatomic particle

  • A loss or gain of electrons changes the atom’s charge, but not its identity

  • Does not significantly affect the mass of an atom

• Atomic notation

  • The atomic mass (protons and neutrons) is written in the top left

  • The atomic number is written in the bottom left

  • Atomic mass and atomic number are measured in AMU

    • AMU - Atomic Mass Unit

    • 1 AMU = 1.66 × 10^-21mg

  • The element symbol is in the centre

• Isotope - Atoms of an element with the same atomic number but different atomic masses

  • Isotopes will have almost identical chemical properties

  • An element will have abundances of each isotope

  • To solve for abundances, multiply each mass by the abundance and add all the factors together

    • Avg AMU = m₁p₁ + m₂p₂ + … m_np_n

  • Half Life - The time it takes for half of an element to decay

  • Radioisotopes - Isotopes that emit energy in the form of alpha, beta, and gamma rays

    • These can be used for practical applications, such as carbon dating, medical imaging, and smoke detectors

• Mass spectrometry

  • A useful technique that can determine the mass of elements, mass of fragments of compounds, and structures of a compound

  • The level that atoms are deflected depends on their mass to charge ratio

  • m/z → Mass to charge ratio

  • A sample will become ionized and is attracted by the magnet

    • Different masses travel at different speeds

    • The detector records the different masses and presents a spectrum

      • The spectrum is presented as a graph, where the mass/charge ratio is on the x-axis, and the abundance is on the y-axis

      • Each peak on the graph is an abundance of an isotope

Lesson 2 - The Quantum Model

• Summary of theories and discoveries

  • Democritus - “atomos,” uncuttable, smallest possible particles

  • Dalton - discovered that matter is made of atoms

    • All matter is composed of extremely small particles called atoms, which cannot be broken into smaller particles, created, or destroyed

    • The atoms of any given element are all identical to each other and different from the atoms of other elements

    • Atoms of different elements combine in specific ratios to form compounds

    • In a chemical reaction, atoms are separated, rearranged, and recombined to form new compounds

  • Perrin - Found that charged particles could be deflected by magnets

  • Thomson - Discovered electrons

    • Found an atom’s mass by deflecting the ray, which depended on the ratio of charge to mass

    • Found particles 2000 times smaller than an atom called electrons

  • Rutherford - Using gold foil experiment, he discovered the nucleus and protons

    • Directed alpha particles at gold foil, found that some bounced back

  • Bohr - discovered that electrons occupy orbits with defined energy

    • Used a spectroscope to identify different levels of energy

    • There were only certain levels of allowed energy

    • If an electron moved in a stationary state, it did not gain or lose energy

    • An electron could move orbits if it absorbed ∆E

  • Balmer - Equation that could determine the wavelengths of the lines in the visible portion of hydrogen’s emission

    • 1/λ = b/hc (1/n₁² - 1/n_h²)

  • De Broglie - Waves of electrons must be stable by being a factor of a whole number

    • Discovered that λ = h/mv

  • Chadwick - Discovered neutrons

  • Planck - Energy could behave like particles called quanta, which could be calculated by E = hv

    • Energy could only be absorbed in whole numbers of quanta

• Spectroscopy

  • How Bohr discovered different energy levels

  • When atoms absorb light, they absorb energy

    • If an electron absorbs the right amount of energy, it will jump to another energy level

    • Once the electron relaxes, it moves back down the energy levels and emits light

    • Metal, when heated, would emits colours, and a prism identified which distinct colours were made

    • A specific wavelength of colour correlates to specific energy levels

    • Therefore, electrons must be in fixed orbits of defined energy

  • Bohr experimented with hydrogen, and calculated the energy levels, which he called quantum (defined quantities)

• Quantum Numbers

  • From Bohr’s calculations, there are 4 quantum numbers

    • Principle Quantum Number “n” - Whole numbers starting at 1 that represent increasing energy levels

      • As energy levels increase, the level is farther away from the nucleus

      • The amount of electrons in an energy level is equal to 2n²

    • Second Quantum Number “l” - The angular momentum number that gives us the shape of an atomic orbital

      • s - spheres - 2 electrons, 1 orbit

        • Located in the alkali/alkaline earth metals

      • p - dumbbell - 6 electrons with 3 orbits

        • Located on the right side of the table

      • d - flower - 10 electrons with 5 orbits

        • Located in the transition metals

      • f - complex - 14 electrons, 7 orbits

        • Located in the lanthanide and actinide series

    • Third Quantum Number “m_l” - The orientation of an atom

    • Fourth Quantum Number “m_s” - The two possible electrons spins, either +1/2 or -1/2

• Electron density

  • Electrons behave as both a wave and a particle

  • Because they are waves, they do not have a fixed position

  • Electron Density - The density around the nucleus that electrons occupy

  • Schrodinger did the calculations for electron density and came up with atomic orbitals

• Filling orbitals

  • Hund’s Rule - When orbitals of equal energy are being filled, the most stable configuration is the one with the maximum number of unpaired electrons with the same spin

  • Pauli’s Exclusion Principle - No 2 electrons in the same atom can be described by the same set of quantum numbers

    • Electrons in the same orbital must have opposite spin

  • Aufbau Principle - Order of filling electrons with the lowest energy first

    • Go from left to right across the periodic table

    • As each orbital shape letter is crossed, fill in the highest possible number of electrons until there are no more electrons to fill in or all orbitals in a level are filled

    • There is a maximum of 2 electrons in n = 1, 8 in n = 2, 18 in n = 3, and 32 in n = 4

• Electron Configuration

  • Write the orbitals that are filled at each level, in the order that they are filled

  • It is written as nl^{(# of electrons)}

  • Noble gas configuration

    • Use the nearest noble gas, and then show the valence electrons that come after

  • For anions, add electrons to the last sublevel

  • For cations, electrons should be removed from the outermost electrons (highest energy level)

Lesson 3 - Periodic Trends

• Development of periodic table

  • First organized as triads

  • Then organized by chemical and physical properties

  • Mendeleev organized columns and rows according the the periodic law

  • Mosely reordered the periodic table based on atomic number

• Elements in the same family have similar configurations, so they have similar properties

• Alkali metals are highly reactive, and alkaline earth metals are slightly less so

• Halogens are very reactive and represent three states of matter

• Noble gases are stable

• Transition elements often have multiple ions

• Periodic Law - When elements are arranged in order of increasing atomic number and therefore, mass, certain sets of properties recur periodically, defined as trends

• The trends are

  • Effective Nuclear Charge (Z_eff) - The net positive charge the outer electrons “feel”

    • How strong the attraction is between the positive nucleus and the negative outer electrons

    • Relies on electrons being attracted to the protons

    • As protons increase, the strength of the nucleus increases

    • Coulomb’s Law - F = (Q₁Q₂)/d² - As the diameter increases, the force of attraction decreases

    • Inner electrons create a shielding effect due to similar charges, repelling the outer electrons

    • This affects the strength of the force on valence electrons and the atom size

  • Atomic/Ionic Radii - The size of the neutral atom

    • Taken from measuring the size of bonded atoms

      • Can be defined by covalent or metallic radius

    • Radius increases as you move down a group because the effective nuclear charge decreases, causing a larger atom

    • Radius decreases as you move right across a period because the electrons are drawn in due to a stronger nucleus

    • Ionic Radius - The size of an ion

      • Cations will be smaller than its atoms because there will be less electrons, and an energy level would be removed

        • The more electrons lost, the smaller it becomes

      • Anions will be larger because it gains electrons, so the effective charge will decrease, and the electron-electron repulsion will be increased

        • The more electrons gained, the larger it becomes

  • Ionization Energy - A measure of how much energy it takes to remove a valence electron from an atom in the gaseous state

    • If more energy is required, there is a higher ionization energy

      • X(g) + energy → X⁺(g) + e^-

      • First ionization energy - the energy required to remove the first valence electron

      • A pair of electrons in an orbital will cause more repulsion, and make it easier to lose an electron, rather than a more stable configuration with one electron in each orbital

    • Ionization energy decreases down a group, as effective energy decreases, so less energy is required to remove an electron

    • Ionization energy increases right across a period, as effective energy increases, so more energy is required to remove an electron

  • Electron Affinity - The energy released when an electron is added

    • X(g) + e^- → X(g)^- + energy

    • It is written in a negative number because energy is released

    • The greater the absolute value of the number is, the more the atom wants the electron

    • Electron affinity decreases down a group, because effective energy is lower, less energy is released when the electron is added

    • Electron affinity increases across a period, because effective energy increases, so it becomes easier to add an electron, so more energy is released

  • Electronegativity - The measure of how strongly an atom pulls electrons towards itself

    • Suggested by Linus Pauling

    • Tug of war concept - the atom with the higher electronegativity will pull the electron towards itself

    • Electronegativity increases across a periodic table, but decreases going down

    • Electronegativity scale predicts whether a bond between two elements will be ionic or covalent using the difference between the electronegativities

      • ∆ EN > 2 = ionic

      • 2 > EN > 1.6 = ionic if it includes metal, covalent if no metal is present

      • 1.6 > EN > 0.5 = polar covalent

        • Not evenly shared

      • EN < 0.5 = nonpolar covalent

        • Very evenly shared

Lesson 4 - Bonding and Structures

• The three types of bonds are covalent, ionic, and metallic

• Ionic Bond - The electrostatic attractive force between the oppositely charged ions produced when a metal atom transfers one or more electrons to a non-metal atom

  • Forms crystal structures called crystal lattices

  • High melting and boiling points because the transfer of electrons makes the bond harder to break

  • Conductive when dissolved in water because anions and cations are charged particles

  • Conductive as molten liquids, but not as solids, because the bonds are compact as solids, meaning ions cannot move to conduct electricity

  • Hard and brittle, because the repulsion between charges causes shattering along specific pathways

  • Lewis notation for ionic compounds

    • Draw each element with its valence electrons

    • Draw the cation without any electrons

    • Draw the anion with a full valence shell

    • Put square brackets around both the cation and anion

    • Include the number of atoms in front of the brackets, and the charge after them

• Covalent bonds

  • Driven by the fact that each atom does not give up electrons easily, so they must share

  • Diatomic elements (HOFBrINCl) must be bonded together when it’s the only element present

  • Low melting and boiling points because of weaker bonds

  • Poor conductor because the atoms are neutral, not ions

  • Solid, liquid, or gas at room temperature because there can be different intermolecular forces involved

  • Can be soft, waxy, flexible, or crystalline because covalent bonds are flexible

  • Lewis structures for covalent bonds

    • Central atoms have the lowest electronegativity

    • Count the total valence electrons

    • Create a bond (2 valence electrons) between the central atom and surrounding atoms

    • Continue to add lone pairs and bonds until all electrons are “satisfied” by the octet rule

      • An atom can break the octet rule to satisfy formal charge if it reaches a higher level

  • Lewis structures for polyatomic ions

    • Central atoms are generally those with lowest electronegativity

    • When counting valence electrons, add an electron if it is an anion, and subtract an electron if it is a cation

    • Create bonds and pairs as if it were covalent

    • Put square brackets around the final bond and include the charge

• Formal Charge - Apparent charges on certain atoms in a Lewis structure that arise when atoms have not contributed equal numbers

  • FC = Valence electrons - Non bonding electrons - Number of bonds

  • Formal charges should be as small as possible

  • The sum of the formal charges in the molecule is equal to the charge on the ion

• Resonance Structure - The possible diagrams to represent the location of delocalized electrons

  • To draw, draw each possible diagram of the molecule

  • Enclose each in square brackets, and draw a double sided arrow between each diagram

  • Include charges if necessary

Lesson 5 - Molecule Polarity

• Dipole - The opposite regions of polarity created when there is a charge differential in a molecule

  • One end has a partial negative charge and the other has a partial positive charge

  • It can be indicated with an arrow that points to the most electronegative atom

  • A nonpolar atom will not have a dipole between atoms

  • A polar covalent will have a dipole between atoms

• Molecular Polarity and Shapes

  • Molecules have a 3D shape

  • The shape is determined by the centre atom and how many atoms are bonded to it

  • If the net of the dipoles move in a direction, then it is a polar molecule

  • If all dipoles cancel each other out (in a symmetrical atom), it is a nonpolar molecule

  • Use a shaded triangle to symbolize a bond coming forward, and a triangle of lines to symbolize a bond moving back

• Electron Group - 1, 2, or 3 bonds, or lone pairs, attached to the central atom

• Electron Geometry - The arrangement of electrons, as described by the bond angle

• Molecular Geometry - The arrangement of the molecule, as described by its shape

• Family of 2 - Linear

  • 2 electron groups on the central atom

  • Bond angle of 180 degrees

  • The atoms on either side of the central atom perfectly repel each other

• Family of 3 - Trigonal Planar

  • 3 electron groups on the central atom

  • Bond angle of 120 degrees

• Family of 4 - Tetrahedral

  • 4 electron groups (4 bonds)

  • Bond angles of 109.5 degrees

  • One bond comes forward, one bond moves back

• Family of 4 - Trigonal Pyramid

  • 4 electron groups (3 bonds, 1 lone pair)

  • Bond angles of 109.5 degrees

  • Polar

  • One bond comes forward, one bond moves back

• Family of 4 - Bent

  • 4 electron groups (2 bonds, 2 lone pairs)

  • Bond angle of 109.5 degrees, except for water, which is 104 degrees

Lesson 6 - Intermolecular Forces

• Properties used to describe compounds

  • Melting Point - The temperature at which a compound changes from a solid to a liquid

  • Boiling Point - The temperature at which a compound changes from a liquid to a gas

  • Solubility - Whether a compound dissolves or not in water

  • Conductivity - The ability to allow electric current to flow (the formation of ions in a solution)

  • Vapour Pressure/Volatility - How easily a compound will vaporize

    • The higher the volatility, the lower the boiling point is

    • Vapour pressure is equivalent to the atmospheric pressure required to break surface tension

• Intermolecular Forces and States of Matter

  • Solid - will vibrate in fixed spots

  • Liquid - slightly further apart than solids

  • Gas - moves very quickly, and can fill a space or be compressed

  • Energy is absorbed to break bonds, and is released to form bonds

  • The amount of energy required to vaporize a molecule depends on their intermolecular forces (IMF)

• Intermolecular Forces - The attraction that keeps molecules together

  • The stronger the force is, the closer the molecules are

  • Ionic bonds are the strongest types of bonds

  • For molecules, the strongest is hydrogen bonding, and the weakest are London Dispersion Forces

• London Dispersion Forces - The temporary dipole created by electron movement, found in all molecules

  • When electrons move to one side, there will be a slight dipole, which will cause neighbouring molecules to form temporary dipoles in reaction

  • When molecular mass increases, London Dispersion Forces increase, as there are more electrons, and a larger radius, meaning they can form a stronger dipole

    • This increases the boiling point, as more energy is required to break it

  • A molecule with branches will have less surface area, and lesser LDF forces

• Dipole-Dipole Forces - The forces between molecules with permanent dipoles, where the negative poles attract to the positive poles

  • Caused by uneven electron sharing

  • They can sometimes repel each other, which causes movement amongst atoms within a substance

• Hydrogen Bonding - A type of dipole-dipole force, which results from a hydrogen atom being bonded to a highly electronegative atom such as nitrogen, oxygen, or fluorine

  • Hydrogen becomes extremely positive and is attracted to the highly negative poles of neighbouring molecules

  • This increases boiling point the most

• Ion-Dipole Force - The attraction between an ion and a polar molecule that causes them to dissolve

  • Occurs when ions are in contact with water, where the ions become more attracted to the polar water molecules than other ions

  • A hydration shell forms around the ion, where the poles of the water molecule orient themselves to the charge of the given ion

• Compounds with similar polarity will mix and dissolve

  • Nonpolar molecules have weak bonds, meaning a nonpolar substance can easily dissolve another nonpolar substance by mixing between its molecules

  • Polar molecules are attractive, meaning a polar molecule can easily attract a separate polar molecule away from its like substance

  • A nonpolar substance can be easily broken up, and a polar molecule will only attract itself, so when mixed together, the polar molecule cannot be dissolved