Unit 1 - Matter and Atomic Structure

Mixtures

• Homogenous: composition is the same in proportion throughout (salt water)

• Heterogenous: composition is different in proportion throughout (stew)

• Elemental: only consisting of one element (O2, Ar)

• Compound: consisting of two or more elements

• Molecules: consisting of two or more atoms (O2, CH4)

• Atoms: consisting of one atom (Au)

Density = m/v

Periodic Table

• Atomic number: top number on the element, equal to the number of protons

• Number of protons: top number on the element, equal to the atomic number

• Electrons: equal to the number of protons if not an ion, otherwise determined by the charge

• Neutrons: atomic mass (bottom of the element) - protons

• Atomic mass: number on the bottom of the element or listed like Titanium-47

Finding the abundance

(6.015122 .07590) + (7.016004 .9241) =6.940 amu

Unit 2 - Light and Electron Configuration 

Light

• Wavelength and frequency are inversely related, the smaller the wavelength, the higher the frequency

• The atom absorbs energy,  causing the electrons to move from a ground state (closer to the nucleus) to an excited state (farther from the nucleus). When the electron returns the atom releases the energy it absorbed in the form of a photon 

Electron Configurations

Begins as 1s22s22p63s23p64s23d104p6… goes to 4th P level then to 3d and from A56 to A57 (6s2 to 4f)

S-2 P-6 D-10 F-14

Oxide ion has -2 (aligning it with the noble gas)

Carbide ion has -4 (aligning it with the noble gas)

Unit 3 - Periodic Table and Trends

Trends

• Atomic Radius: ⬇⬅ (Francium biggest) Atomic radius decreases across a period (row) since effective charge increases along a period, causing the attraction to be higher and making the electron cloud smaller. 

• Ionic Radius: ⬇⬅ (Francium biggest)

• Ionization Energy/Energy required to remove an electron: ⬆➡ (Helium highest) Ionization energy is higher across a period since the effective nuclear charge is higher, making removing it more difficult. Ionization decreases down a group (is easier to remove) since energy levels increase and therefore the balance electrons are further apart and less attracted to the nucleus. 

• Electronegativity/Ability to attract electrons: ⬆➡ (Fluorine highest) Electronegativity increases across a period (row) since the effective nuclear charge increases across a row. It decreases down a group since energy levels increase down a group and make valence electrons further from the nucleus, therefore making attraction more difficult. 

Metals vs Nonmetals

• Metals: malleable, high melting point, low electronegativities, low ionization energy, metals lose electrons

• Non-metals: non-malleable, varying melting points, higher electronegativities/ionization energy, non-metals gain electrons

Unit 4 - Bonding

Lewis Structures

Bonds

• Ionic: Metal and Non-metal, electrons are transferred from metal to the non-metal

• Soluble and conductive

• Polar Covalent: Non-metal and Non-metal with different electronegativities, electrons are shared unequally

• Soluble but not conductive

• Non-polar Covalent: Non-metal and Non-metal with similar or the same electronegativity, electrons are shared equally

• All no: Doesn’t dissolve in water or conduct electricity. Not malleable.

• Metallic: Metal and Metal: None want the electrons, creating a sea of electrons creating an electrical current and conducting electricity

• Conducts electricity in bulk phase, is malleable. 

Polyatomic ions can have both covalent and ionic bonding. Ionic between the counterion and polyatomic ion, covalent between the atoms within the polyatomic ions. 

Unit 5 - Nomenclature

Naming

• Transition metal: include the charge in name

• Non-metals: include a prefix, always on second, on first if not “one”

• The second element gets an ide (not polyatomic ion)

• Polyatomic ions leave as is

Unit 6 - Moles and Empirical Molecular Formulas

Empirical/Molecular Formulas

1. Put percent as a mass

2. Convert mass to moles 

3. Divide by the smallest mole number

4. Multiply until whole number

5. To make it molecular take the empirical formula and add up the mass and then if not the mass multiply until it is 

Unit 7 - Balancing Reactions and Predicting Products

• Combustion produces CO2 and H2O

• Double Replacement reactions occur if a precipitate is formed

• For Single Replacement reactions to occur, the metal must be higher on the activity series than the one it is trying to replace

Unit 8  - Stoichiometry

• Use a molar ratio

• Whichever makes the least product (solve for the same one) is the limiting reactant

• Percent yield = actual yield/theoretical yield x 100

  • When calculating what will be produced (product), multiply

  • When calculating what needs to occur for something to be produced (reactant), divide

Unit 9  - Gases

• PV=nRT

• PV/NT=PV/NT

• Divide = direct, Multiply = Inverse

• Add 273 to Celsius

Unit 10  - Solutions

• Molarity = mol of solute/liters of solution

  • Ensure convert to liters (and grams to moles)

• MV=MV

• Saturated = can’t add any more

• Unsaturated = can add more

Unit 11 - Acids and Bases

Acids

• Source of H+ (ex: HCl, H2SO4

• Ionic bonding (soluble and conductive, categorized as an electrolyte)

• pH under 7

• Sour

• Turns clear

• Neutralize with baking soda (a weak base)

Bases

• Source of OH- (NaOH, KOH)

• Ionic bonding (soluble and conductive, categorized as an electrolyte)

• pH over 7

• Bitter

• Turns pink

• Neutralize with vinegar (a weak acid)

• A specific kind of double replacement reaction

• Acid + Base -> Salt (Conjugate Base) + Water (Conjugate Acid)

• In a Bronsted-Lowry reaction, protons are transferred

  • Bases accept H+ ions

  • Acids donate H+ ions

• Strong acids completely dissociate, weak acids do not

• If it is an ionic bond (metal and non-metal) or contains a polyatomic ion it is an electrolyte