Honors Chemistry Mid-Term

→Elements/Polyatomics:

Know 50 elements (I ain’t gonna write it)

Polys:

-Ammonium NH4 (+)

-Sulfite SO3 (2-)

-Sulfate SO4 (2-)

-Hydrogen Sulfate HSO4 (-)

-Nitrite NO2 (-)

-Nitrate NO3 (-)

-Cyanide CN (-)

-Perchlorate ClO4 (-)

-Chlorate ClO3 (-)

-Chlorite ClO2 (-)

-Acetate C2H3O2 (-)

-Hypochlorite ClO (-)

-Chromate CrO4 (-)

-Dichromate Cr2O7 (2-)

-Permanganate MnO4 (-)

-Thiocyanate NCS (-)

-Mercury (I) Hg2 (2+)

-Oxalate C2O4 (2-)

-Peroxide O2 (2-)

-Phosphate PO4 (3-)

-Hydrogen Phosphate HPO4 (2-)

-Dihydrogen Phosphate H2PO4 (-)

-Hydroxide OH (-)

-Carbonate CO3 (2-)

-Hydrogen Carbonate HCO3 (-)

→Scientific Method:

1. Observation - what do you see?

2. Hypothesis - an educated guess based on observation

3. Experiment - testing hypothesis

4. Conclusion - results → was your hypothesis supported or not supported? If not, change it.

→Measurement:

• Measurement → quantitative observations (unit, number)

• Exact numbers → #s you can count (have an infinite number of sig figs)

• Scientific Notation → same # of sig figs in answer (6.02 × 10 23 = 3 SFs)

Important: Measure one place beyond what is marked on the instrument you are using

• Rules for Sig Figs-

  1. Any non-zero # is significant (714, 3 sig figs)

  2. Any zero between two non-zero # is significant (107, 3 sig figs)

  3. Any zero before the first non-zero # is not significant (0.0238, 3 sig figs)

  4. Any zero holding place is not significant (100, 1 sig fig VS 100. 3 sig figs)

  5. Any zero after a # & after a decimal is significant (0.7030, 4 sig figs)

• SF Rules for Multiplication/Division

  • the answer has the same # of SFs as the quantity in the problem with the least # of SFs

    • EX: 123 x .27 = 33 (27 has only 2 SFs), 1463/279 = 5.88 (279 has only 3 SFs)

• SF Rules for Addition/Subtraction

  • the answer must stop at the same place as the measurement with the fewest/smallest decimal places

    • EX: 16.1 + 2.37 = 18.5 (16.2 is only one decimal place), 0.6785 - 0.11 = .57 (0.11 is only 2 decimal places)

• Dimensional Analysis

  • # & unit (must have)

  • units in first, then #s

  • always convert to the base unit

  • MULTIPLY, MULTIPLY, MULTIPLY, then, DIVIDE, DIVIDE, DIVIDE

  • Must Know: 5,280ft = 1 mile, 2.54cm = 1 in, 1cm3 = 1 mL, 454g = 1lb

    • EX: Convert 1.728 km to nm- (1.728km)(10³ m . 1 km)(1 nm / 10⁻⁹ m) = 1.728 × 10¹² nm

  • Conversions- going from km per day to mi per hr, go for top first, then bottom

  • Conversions w/ Squares & Cubes- square & cube the # & unit

    • EX: 15.0 yd² to ft²- (15.0 yd2)(3ft/1yd)² = 135 ft²

→ Chemistry & Matter:

• Chem → the study of matter & change of matter

• Matter → anything that has mass & occupies space

  • Solids- definite shape, definite volume

  • Liquids- definite volume, indefinite shape

  • Gases- indefinite shape, indefinite volume

• Mass → amount of matter

• Density = mass/vol

• Viscosity → resistance to flow

  • EX: Honey is high, Water is low

• Matter → pure substances (elements & compound) & mixtures (homogeneous & heterogeneous)

  • Homogeneous- appear the same throughout, uniform characteristics

    • EX: Air, Soda, Gasoline, Saltwater

  • Heterogeneous- don’t appear the same throughout, different characteristics

    • EX: Granite, Concrete, Sand

• Separating Matter

  • Filtration- separate the components of a mixture

  • Chromatography- passing it in solution to vapor

  • Distillation- separate mixtures that are comprised of 2 or more pure liquids

• Properties

  • Intensive- do not depend on the amount of material

    • EX: Color, Density, Temperature

  • Extensive- does depend on the amount of material

    • EX: Mass, Volume, Energy

• Classifying Matter

  • Physical Property- any property of a physical system that is measurable

    • EX: Color, length

  • Chemical Property- a property that is measured or observed following a chemical change to a substance

    • EX: Toxicity, flammability

  • Physical Change- changes affecting the form of a chemical substance, but not its chemical composition

    • EX: Crumbling paper, ice melting

  • Chemical Change- a change of materials into another, new materials with different properties, and one or more than one new substances are formed

    • EX: Digestion of food, baking a cake

→Atomic Structure:

• Scientists Importance

  • Dalton- 1st model, blue sphere, fixed ratios form compounds

  • Thomson- discovered electrons; Plum Pudding Experiment; mass-to-charge ratio

  • Milikan- Oil-Drop Experiment, also discovered electrons

  • Rutherford- atom is mostly empty space, and mass is in the positively charged nucleus; Gold Foil Experiment; discovered protons

  • Chadwick- discovered neutrons

  • Bohr- orbit model

  • Schrodinger- orbital model, most current

• Subatomic Particles

  • # of protons determines what the element is (Periodic Table)

  • Neutrons → inside nucleus; positive charge

  • Electrons → outside nucleus; negative charge

  • Protons → inside nucleus; no charge

  • Isotopes → same element, different atomic masses

  • Ions → charged particles

    • Anion → negatively charged ion

    • Cation → positively charged ion

  • ATOMIC MASS = PROTONS + NEUTRONS

    • EX: Look in notebook

• Isotopes

  • Def: atoms of the same element, with different # of neutrons/atomic mass

    • amu- atomic mass unit (1/12 of the mass of a carbon-12 atom)

  • Mass Spectrometer → abundances of isotopes in nature

    • EX: Look in notebook

  • Uses of Isotopes

    • fuel in a nuclear reactor

    • treatment for cancer

    • smoke detector

• Electrons

  • Ways to Represent Electrons (Look in notebook for help)

    1. Electron Configuration - s left, d middle, p right, f bottom

       • Important: when we remove electrons, remove them from the highest energy level

    2. Orbital Diagram

       • Singly 1st (Hund’s Rule)

    3. Lewis Dot Structure

       • elements symbol & # of valence electrons

• Energy Problems (Look in notebook for problems)

  • Need to know:

    • E = hv

    • C = lv

    • h = 6.626 × 10⁻³⁴ Jxsec

    • C = 3.00 × 10⁸ m/sec

→Naming Compounds & Writing Formulas:

• Periodic Table

  • Horizontal Rows- periods

  • Vertical Columns- families or groups

• Groups

  • Group 1- Alkali Metals (+1)

  • Group 2- Alkaline Earth Metals (+2)

  • Group 3-13- Transition Metals

    • Exceptions in- Ag (+1), Zn (+2), Cd (+2)

    • Exceptions out- Sn, Pb

  • Group 13- Boron Group, Al (+3)

  • Group 14- Carbon Group, no charge

  • Group 15- Nitrogen Group (-3)

  • Group 16- Oxygen Group (-2)

  • Group 17- Halogens (-1)

  • Group 18- Inert/Noble Gases, no charge

• Blackline

  • left = non-metal

  • right = metal

• Does the Compound Contain a Metal?

  • Yes = Ionic Compound

  • No = Molecular Compound

• Ionic → name of metal, root of anion, -ide

  • Charges add up to equal zero when it’s named to compound

    • EX: NaCl → Sodium Chloride

• Molecular → prefix + 1st element name, prefix, root, -ide

  • Name to formula → name what it’s telling you to write

    • EX: CO2 → Carbon Dioxide

• Transition Metals

  • indicate charge using Roman numerals (in name, not formula)

  • Roman numerals is the charge, not the number of that metal

    • EX: CuCl → Copper (I) Chloride

• Polyatomic Ions - Cations first, anions second

  • charges add up to equal zero

  • if more than one PA is needed, use ( )

    • EX: Ca(OH)2 → Calcium Hydroxide

• Prefixes

  • one = mono

  • two = di

  • three = tri

  • four = tetra

  • five = Penta

  • six = hexa

  • seven = hepta

  • eight = octa

  • nine = nona

  • ten = deca

→Chemical Reactions:

• Chemical Reaction → when matter combines or breaks apart to produce new kinds of matter with different properties

• Four Evidences of a Chemical Reaction:

  • Color Change

  • Release of a gas

  • Precipitate → solid that forms from solutions being combined

  • Change in heat or light

• Reactants → starting substances (written on the left)

• Products → the finishing substances (written on the right)

• Important Note: when writing and balancing, the equation must have the same number and type of each element on both sides of the chemical reaction

• Symbols:

  • ¨+¨ → add to; reacts with; mixed with

  • ¨→¨ → yields; produces; forms

  • (s) → solid; (l) → liquid; (g) → gas; (aq) → aqueous

  • arrow with a triangle above → Heat was added

  • arrow with a formula above → catalyst (a substance that speeds up the rate of a reaction without being used up in the reaction by lowering activation energy)

  • two half arrows facing opposite directions → reaction of reversible

  • Î → gas was released

  • down arrow → precipitate was formed

  • NR → no reaction took place

• Diatomics → Hydrogen, Nitrogen, Oxygen, Fluorine, Chlorine, Bromine, Iodine

• Types of Reactions:

  • Combination Reaction: element + element → compound

  • Decomposition Reaction: 1 reactant → 2 or more products

  • Single Replacement Reaction: element + compound → compound + element

    • Activity Series Chart

    • If the element is higher than the cation, then it will replace it. If not, no reaction

  • Double Replacement Reaction: 2 reactants + 2 compounds → 2 products + 2 compounds (precipitation reaction)

  • Combustion Reaction: same products CH4 + 2O2 → CO2 + 2H2O

• Acids:

  • Hydrochloric → HCl

  • Sulfuric → H2SO4

  • Phosphoric → H3PO4

  • Nitric → NHO3

• Write & Balance Reactions:

  • Write the formula for reactants to the left, the formula for product to the right

  • Balance equation using coefficients (number in front of formula)

    • When balancing, do hydrogen and oxygen last

• Other Types of Reactions:

  • Exothermic: exo- → out; releasing heat; reactions have more energy than products (combustion reaction)

  • Endothermic: endo- → takes in; absorbs energy; products have more energy than reactions

→Mole & Stoichiometry:

• Stoich → the study of the amount of reactants & products consumed & produced in chemical reactions

• The Mole (mol) → the # equal to the # of carbon atoms in exactly 12 grams of our carbon-12

• Mole = 6.022 × 10²³ (avogadro’s #)

• Molar Mass

  • the mass in grams of one mole of a substance

  • the mass of one mole of an element is equal to its atomic mass in grams

  • Molar mass = gram/mole

• Molar Mass for Compounds

  • EX: H2O

    • H: 2 × 1.01 = 2.02

    • O: 1 × 16.00 = 16.00

    • combined = 18.02 g/mol

• Mole Conversions

  • EX: 12.8g H2O to moles

    • (12.8g H2O) (1 mol H2O/18.02g H2O) = .71 mol H2O

• Empirical & Molecular Formulas

  • Empirical → smallest whole # ratio of atoms in a compound

  • Molecular → actual formula of a compound

  • Steps:

    1. figure out # the grams of each element in the compound

    2. calculate # the moles in the n compound

    3. divide by smallest # of moles in Step 2

    4. write empirical formula (multiply as needed)

    5. calculate the molar mass of the empirical formula

    6. take the molar mass of the molecular formula (given), divide it by the EF molar mass

    7. multiply the subscripts of the empirical formula to # in Step 6 (whole #)

• Stoichiometry

  • NTK: ( actual yield / theoretical yield ) x 100 = percentage yield

  • Look at notebook for problems

• Lab

  • Percent composition = (mass of element (g) / total mass (g)) x 100

  • Percent error = ( | theoretical value - experimental value | / theoretical value ) x 100