Mid-Term Exam Review

Lesson 1 – Average Atomic Mass

Mass Number = Atomic Number (Represents number of Protons) + number of neutrons

Abundance of Each Isotope

o   u = Mass 1*Abundance 1(in decimal) + Mass 2*Abundance 2 …

o   Avgamu = m₁p₁+m₂p₂…

o   Must write out for full communication Marks.

• Must multiple the relative abundances of the element’s isotopes by their atomic masses then summing the products

• the average atomic mass on the periodic table is the average of all the masses of all the isotopes

  Isotopes

o   Most abundant when the -n number is closes to the one on the periodic table.

o   To find the abundance from the mass must use:

o   Uses the thing on the periodic table form Avg amu.

o   There are 2 equations one is always about p1+p2=1 as they must equal 100%

o   Rearrange for the x and substitute.

Lesson 2 – The Quantum Model

“n”

o   Whole number starting at 1 only up to 7.

o   Goes 1=2, 2=8, 3=18

o   Electrons, waves, and particles

o   Never in a fixed position

o   There is probability density of where the electrons could be at that moment but is always moving around (ELECTRON DENSITY) → Never actually know where an electron is at any given moment.

o   Schrödinger created this and made ATOMIC ORBITALS

Atomic Orbital Shapes

o   S = Spheres = 2 electrons (1 Orbital)

o   P = dumbbells = 6 electrons (3 Orbitals)

o   D = flower-like = 10 electrons (5 Orbitals)

o   F = complicated flowers = 14 electrons (7 Orbitals)

Rules of Filling

o   Hund’s

o   Half fill first then full fill

o   Pauli’s Exclusion Principle

o   No electron can have same set of 4 quantum numbers (ONE SPINS ONE WAY AND THE OTHER SPINS THE OTHER WAY)

o   Aufbau Principle          

o   When doing the filling must make sure that after 56 (Ba) must go to the 57 or La so go down and that starts at 4f and then only has also 5s

Polar VS None Polar Electronegativity

• under 0.5 → non-polar covalent

• 0.5-1.6 → polar covalent

• 1.6-2.0 → polar covalent if 2 non-metals or ionic is 1 non-metal and 1 metal

• over 2 → Ionic

• could also be angles with VSEPR (already in note with the bond angles)

Periodic Trends

o   Periodic Law = Something that repeats and follows a pattern

Effective Nuclear Charge

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

o   As you increase the number of protons you increase the force of attraction

o   As you increase the number of inner electrons you decrease the force of attraction using the SHEILDING EFFECT

Atomic Radius

o   The size of the NEUTRAL ATOM

Ionic Radius

o   The size of an ION (IN COMPARISON TO THE NEURTAL ATOM OR OTHER ATOMS)

o   Cation: become smaller as loses electrons POSITIVE CHARGE

o   Anion: become bigger as there is an addition of electrons NEGATIVE CHARGE

Ionization Energy

o   How much energy it takes to remove ONE VALENCE ELECTRON

o   More energy = higher ionization energy

o   Atom_(g) + energy = ion⁺+ e^-

Electron Affinity

o   Opposite of Ionization Electron

o   How much energy RELEASED with one electron is added.

o   Must be released to be able to add the electron to decrease the stress on the atom.

o   This is a NEGATIVE number.

o   Greater negative = great want for an electron

o   i.e. Florine has the highest electron affinity, so they are wanting the electron the most.

Electronegativity

o   An atom’s tendency and want to attract the electron shared in a bond or strip in an ionic compound of its valence electrons.

Lesson 4 – Bonding and Structures

Structures of Ionic Bonds and Properties

o   Crystal LATTICE STRUCTURES

o   High boiling and melting point (Strong Bonds)

o   Conducts electricity when dissolved in water (IONS)

o   Conducts electricity as molten liquids but not as a solid (DOES NOT MOVE SO CAN’T CONDUCT)

o   Hard and brittle (repulsion between charges)

Structure of Molecule and Properties

o   Low melting and boiling point (weaking bonds)

o   Poor conductivity (Neutral no ions)

o   Solid, liquid, gas at Room temperature (Depends on the intermolecular forces involved)

o   Usually soft and waxy, flexible (BONDS ARE FLEXIBLE)

o   Can be crystalline but that is an exception.

o   Electronegativity

o   LINUS PAULING FOR IMPORTANT PROPOSION

o   Less than .5 – nonpolar

o   .5-1.6 -polar

o   More then 2 – ionic

Lesson 5 – Molecule Polarity

Dipoles

o   This is a vector.

o   is created when their ware regions of negative and Positive charge.

o   NOT FORMAL CHARGE but when added all together can find formal charge, must have a full symmetry of all outside elements in the molecule to POSSIBLY either a polar or non-polar element.

Shapes of Molecules with Molecule Polarity

o   Electron Groups – one double, single, or triple bond or lone pair

o   Electron Geometry – arrangement of electrons (EG) → number of electron groups

o   Molecular Geometry – shape of molecule (ONLY LOOK AT THE CENTER OF THE ATOM) à how they are arranged

Family of 2 – Linear

o   Has Two Electron Groups

o   They are as repelled as possible, so they are balances.

o   Non-polar as dipoles cancel each other out, if the same element

• Polar if the elements are of different electron negativity

   Family of 3 – Trigonal Planar

o   Has 3 electron Groups

o   If symmetrical then non-polar

o   Some middle elements are happy with 6 electrons

Family of 4

   Tetrahedral

o   4 electron groups (4 bonds)

o   One forward, one left, one right, one back

o   Non-polar, unless onto all the same then polar

    Tetrahedral, Trigonal Pyramidal

o   4 electron groups, one lone pair, 3 bonds

o   Polar

o   Side, forward, back

o   109.5

  Tetrahedral, Bent

o   4 electron Groups, 2 orbitals, 2 bonds

o   Polar

o   one left, one right, orbitals on top

o   WATER: BOND ANGLE IS 104.5

o   NORMAL: 109.5

Lesson 6 – Intermolecular Forces (IMFs)

Attraction that keeps molecules together, stronger forces the closer the molecules.

Properties of Compounds

   o     Melting Point: Solid to liquid

   o     Boiling Point: liquid to gas

   o     Solubility: can it dissolve into water (LIKE DISSOLVES LIKE) (POLAR DISSOLVES POLAR)

   o     Conductivity: Allows electric Current (Only with Ions)

   o     Vapour Pressure or Volatility: balance between particles in that has and water phases. Higher volatility is lower boiling point.

London Dispersion Force (LDF)

   o     ALL COMPOUNDS HAVE THIS

   o     This is when there is a temporary dipole as the electrons are always moving for cloud to cloud.

   o     This attracts each other in a compound.

   o     Typically, the weakest IMF unless it has an extremely high molecular mass which would lead to a high LDF.

o   Molecular mass directly affects LDF as they are more loosely held, also this would increase the surface area which give more space for the electrons to hop around.

Dipole-Dipole Force

   o     Always have a permanent dipole with polar molecules

   o     Electrons are always being pulled in one direct or another so there will always be a charged side that will be attracted to it à forms a link of SAME COMPOUND TOGETHER

   o     Like electrostatic forces between ions but you don’t have fully charged atoms, just an uneven electron sharing.

Hydrogen Bonding

   o     Special type of dipole force between hydrogen and Nitrogen, Oxygen, or Florine

   o     Hydrogen is strongly negatively charge; this is the strongest IMF.

When there is an addition of molecular mass, dipole-dipole, or hydrogen bonding the melting and boiling point will increase as the forces become stronger.

 Organic Naming

Unit 2 Reactions

KNOW SYNTHESIS, DECOMPOSTION, SINGLE DISPLACEMENT, DOUBLE DISPLACEMENT

Combustion Reactions

   o     C_nH_m + O_{2 →} CO_2(g) + H₂O_(g) → Complete

   o     C_nH_m + O_{2 →} CO_2(g) + H₂O_(g) + C_(s) + CO_(g)

   o     Aqueous solutions dissolved in water à the molecule will disassociate want to become the ions.

Neutralization Reactions

   o     Acid (H⁺) + Base (OH^-) → Ionic salt + water (LIQUID THAT COULD BE THE PRODUCE FOR NET IONIC EQUATIONS)

   o     Can also be a gas and make CO_2(g) and H₂O_(l)  Which is with sulfides and Carbonates

Particulate Diagrams

   o     Show how particles are behaving in solutions.

   o     Circles are one particle and are annotated to become molecules.

• FOR AQUEOUS ALL PARTS MUST BE SEPERATED YET AROUND ONE ANOTHER

  

Net Ionic Equations

1.      Balanced chemical equations for the reaction.

2.      Write out total ionic equation à all subscripts become number in the front à must add charge.

a.      Leave the solid compound as is à DON’T SPLIT IT

3.      Cancel out any ions that do not appear in the solid compound à THEY STAY THE SAME STATE

4.      Rewrite final net ionic equation à WITH IONS THAT MAKE THE SOLID

Example with steps:

Redox

   o     Oxidation = loss of election = reducing agent

   o     Reduction = gain of election = oxidizing agent

Oxidation number

o   Changes of the atom would have if electrons were not started by transferred.

o   SIGN BEFORE

   o     RULES FOR OXIDATION NUMBERS

o   Elemental form: N.O. = 0

o   Monoatomic form: N.O. = ionic charge

o   Sum of N.O. values for a compound in a molecule =0

o   Polyatomic: N.O. total = ions charge

Common/specific values for groups

o   Oxygen: -2, unless peroxide (Y_xO_x) then oxygen is a charge of -1

o   Hydrogen: +1 when nonmetals, -1 with metals

o   Group 17: -1 in combinations with metals and nonmetals (WITH OXYGEN THE HALOGEN TAKES THE + UNLESS ITS FLORINE THE -)

o   Group 2: +2 in all compounds

o   Group 1: +1 in all compounds

  Reactions

1.      Assign all oxidation numbers in elements.

2.      Determine if any changes in oxidation numbers have occurred.

a.      Increase or decrease.

b.      Use the chart!!!

3.      BOTH INCREASE AND DECREASE NEED TO HAPPEN IF A REDOX HAPPENS

   o     Half reactions

o   Oxidized is on the right.

o   Reduced is on the left.

o   When calculating how many electrons you are working with, they must equal each other in both the reduction and oxidation.

                     o     The electron you are working with is equal to the coefficient multiply by the coefficient multiplied by the oxidation number.

                     o     When writing keep the subscripts and the coefficients in the normal place and do not multiple

Nuclear Reactions

Nuclear Half Life

   o     The time it takes for ½ of the quantity of the substance to decay.

   o     The radioactive isotopes emit radiation, so they decay themselves.

   o     TO CALCUATE: A_F = A_I(½)^t/p à just use Log to find the t or p

   o     Uses in reality: x-rays, carbon dating (C-14) , to kill cancer with radiation,

Nuclear Reactions

   o     Fusion: lighter nuclei combine into a heavier one

o   Energy is released.

o   The sun

   o     Fission: heavier nucleus splits into lighter ones

o   Energy is released.

o   Nuclear power plants and atomic bombs

Types:

 Unit 3 - Stichometry

Significant Digits

• non-zeros are significant

• Sandwich zeros and trailing zeros AFTER A DECIMAL are significant

• leading is NEVER significant

Multiplication and Division

• just least number of significant figures

Addition and Subtraction

• Least number of decimal places (if no decimal then just place value/# of sig figs)

Mixed

• Use BEDMAS

• Round everything then use it on the calculator but write rounded

Molar Relationship

• Just g/mol^-1 so the number of the element times the molecule mass

Percent Composition

• Tells you how much of the mass is that element in percent

• Formula: MEMORIZE

% of Element = Molar Mass (Periodic Table, in g*mol^-1) * # of atoms of element in chemical Formula DIVIDED BY the Molar mass of the compound (The # of Molar masses added)

Empirical Formula

Simplified/most reduced ratio → proportions of each element in a compound

Steps:

1. Assume you have 100g (so 1% = 1g) of the sample (if not give mass in grams, then just use that)

2. Covert % of element to mass (grams)

3. Convert Mass to Moles

   • Do this by multiple the full number by the reciprocal of the period table number 1/x or just divide by the mass (in g/mol)

4. Determine the WHOLE-NUMBER ratio of moles by dividing by the smallest number of moles (smallest of the number of element you have) → if not in whole numbers just multiple by the denominator of the fraction its in BY ALL OF THEM TO STILL GET THE RATIO

5. Then just write the ratio in a therefore statement

Molecular Formula from Empirical Formula

The actual ration of elements in a given substance

USE THE STEPS ABOVE TO FIND THE EMPRICAL FORMULA THEM USES FORMULA BELOW TO FIND THE FINAL

NEED TO KNOW Empirical formula and molar mass

Formula: MEMORIZE

Molecular Formula Mass DIVIDED BY Empirical Formula = RATIO (what you have to multiple by each subscript to get the formula

Empirical Formula → the Mass on the period table multiplied by the number in each \

MUST SHOW ALL OF THE MUTIPLCATION F THE RATIO AND THE FORMULA