Chemistry Final Review

Name:______Chemistry Final Review

1) Differentiate between solids, liquids, and gases: plasma liquid Solid gas fixed shape (yes/no) fixed volume (yes/no) flows (yes/no) compressible (yes/no) Order from least(4) to greatest(1) in the following categories: particle proximity molecular motion Arrangement average kinetic energy intermolecular attraction 2) Temperature is the measure of a substance’s average [kinetic/potential] energy that increases as its temperature [increases/decreases]. 3) Physical property: ______4) Chemical property: ______5) Classify as physical (p) or chemical(c) property of a substance: boiling point __, rusts __, color __, reacts w/acid __, melting point __, density __, decomposes __, texture __, burns __ 6) Physical change: ______7) Chemical change/chemical reaction: ______8) 6 signs a chemical change:______, ______, ______, ______,______, ______9) Classify as physical (p) or chemical (c) change: burning __, dissolving __, rusting __, melting __, evaporating __, bleaching __, subliming __, baking __, decomposing __, boiling __, rotting __, cutting __ 10) Indicate the state changes: solidliquid ______; liquidgas ______; solidgas ______gasliquid ______; liquidsolid ______; gassolid ______11) Melting, boiling, and sublimation are the state changes that result from an [increase/decrease] in energy. 12) Freezing, condensing, and deposition are state changes that result from [increasing/decreasing] energy. 13) Atomic Theory states that ______14) What 2 principles of Dalton’s Atomic Theory changed? How? i. ______ii. ______Name:______15) Thompson discovered ______w/the ______experiment. His ______model stated negatively charged e-s are embedded in positively charged matter/atom. 16) Rutherford discovered the ______w/the ______experiment. In his model, negatively charged e-s are scattered around the small, dense, positively charged nucleus that contains most of an atom’s ______. 17) Bohr’s Model- electrons can only be in ______around the nucleus. Like ______on a ladder. 18) Quantum Theory takes into account both ______and ______properties of electrons. 19) The number of ______and ______in the nucleus is mass number; Atomic number, which identifies the element, is the number of ______. 20) The modern periodic table is based on atomic ______. Mendeleev’s PT was based on atomic ______. 21) Isotopes are atoms of the same element with different numbers of ______and different ______. 22) List the subatomic particles (p+, n0, e-) in order of decreasing mass.____, _____, _____

23) Complete subatomic particle structure charts: subatomic Charge Mass # Mass Location in the atom (nucleus, outside the nucleus) particle electron proton neutron Hyphen isotopic symbol Mass# (p++n0) Atomic # (p+) p+ n0 notation Boron-12 81 1- 35N 16 18 24) ______:all frequencies(ν) or wavelengths(λ) of electromagnetic radiation. 25) Frequency(ν) + wavelength(λ) are ______related; if frequency(ν) decreases, wavelength(λ) ______. 26) Show substitution and rearranged equation solving for the unknown variable. c = νλ, c = 3.00x108m/s, c = speed of light in a vacuum, 1 hertz (Hz) = or 1s-1 ex) Calculate wavelength(λ) of microwaves used to cook food with a frequency(ν) of 2.76x1010Hz. 3.00x10 8 m/s=(2.76x10 10 Hz) λ  λ =(3.00x10 8 m/s)÷( 2.76x10 10 Hz)= 1.09x10-2m a) What’s the wavelength of violet light with frequency of 2.59x1014 Hz? ______=1.16x10-6m Name:______b) What’s the frequency of light reflected from a green leaf with a wavelength of 8.95x10- 7m? ______=3.35x1014Hz 27) Determine energy (in joules) of a photon w/ frequency of 5.35x1017Hz. Show substitution. E=hν, E=energy, h(Planck’s constant) =6.626x10-34J·s ______=3.54x10-16J 28) Photoelectric effect: ______29) Ground state: ______30) When excited an e- has [gained/lost] energy and has a [higher/lower] potential energy than in ground state. 31) In dropping from a higher energy state to a lower energy state, an electron emits ______as ______. 32) ______: specific frequencies of light (spectrum of a few colors) emitted from elements (each element is unique) resulting when light passes through a prism and separates. 33) List the 4 quantum numbers, what they mean, and fill in the chart. a) P ______: ______b) A______: ______c) M______: ______d) S______: ______sublevel shape # orbitals or orientations about the nucleus Max #e-s in sublevel s p d f 34) Pauli exclusion principle: ______35) Aufbau principle:______a) The energy levels in increasing order: 1s2, ___ , ___ , ___ , ___ , ___ , ___ , ___ , ___ , ___ , ___ , 6s2 36) Hund’s rule: ______37) Electron configuration (superscripts) for Co: ______a) How many electrons are in cobalt’s electron configuration? ___ Name:______38) Orbital () notation for N: ______a) # of electrons in N configuration? __ 39) Which element has an electron configuration that ends with 4p2? ___; ends with 3s2? ___; 5d9?____ 40) What element has a noble gas configuration of [Kr]5s24d105p4? ___; [Xe]6s1?___; [Ar]4s23d3?___ 41) Noble gas configuration of exceptions *[Ar]4s13d10? ___; *[Ar]4s13d5?____ 42) How many electrons are required to fill the 1st energy level (period)? __; 2ndenergy level? __;4th?___ 43) Valence electrons are electrons in the ____ and ____ orbitals of the outermost energy level. 44) When an atom has a full outer energy level of 8 electrons it is [less/more] stable than when it doesn’t. 45) Elements in the same ______, column, have the same number of valence electrons. They have similar ______and ______properties and therefore, behave ______. 46) Elements in the same ______, row, have the same principle quantum number, thus have the _____ number of occupied energy levels. 47) How many partially filled orbitals are in the following atoms: P______; S______; Li______; Cl______48) Most elements on the periodic table are [solid/liquid/gas/plasma]. 49) What element is in period 3, group 2? __; group 14, period 4? __; P 3, G 18?__; G6 P5? __ 50) Identify each group of the Periodic Table and give 1 characteristic of each:

A) E) ______;______;______B) F) ______;______;______C) G) ______;______;______D) H) ______;______;______Name:______I) J) ______;______;______51) Underline + state how many significant figures: (leading zeroes aren’t significant, trailing are w decimal) a) 10000 b) 1.00

c) 0.010

d)0.0110 e) 1.000x109

f) 1x10-7 g) 1.0000

h) 0.0001

i) 10.100

j) 1100 k) 1.0x1020 l) 1.00x104

52)Round ↓ to 1 sig figs Round ↓ to 2 sig fig Round ↓ to 3 sig figs Round ↓ to 4 sig figs 0.034250.03 0.034520.035 34.56734.6 912678912700 625712 1.0784 678341 2399.9 53)scientific notation↔standard notation ex)1000↔1x103; 0.001↔1x10-3 a)0.000375= b) 8.68x10-3= c) 7000= 54)Density (D=m/v): Show substitution, work, intermediate answer, final answer to correct # of sig figs. a)What is the mass of 100.0cm3 of gold if the density is 19.32 g/cm3? 1932

b)What substance has a mass of 55.0g and occupies 6.14cm3? Fe=7.87g/cm3 Cu=8.96g/cm3 Pb=11.36g/cm

55) In an experiment, a/n ______variable (cause) is a variable that is varied or manipulated, a/n ______variable (effect) is the response that is measured, and a/n ______is a constant variable. 56) In an experiment designed to test “How does concentration effect reaction rates?“ independent variable ______dependent variable ______control ______. Name:______57) Ion formed (symbol+charge):radium____;aluminum____;bromine___;phosphorous___;sodium____;seleniu m___ 58) Name transition metal

ion:Cr2O7______;CuO______;FePO4______;AgI2______59) Determine the number of valence electrons for: Ar=__; Be=__; P=__; Al=__; K=__; F=_; Si=_; S=_; He=_

60) Identify each compound as Ionic (I), Covalent(C), or Acid(A) and give the name or formula: Type Formula Name

Type Formula Name a) I

nitrogen triiodide c)

aluminum chloride e)

hexacarbon decahydride g)

Na2S Name:______h)

vanadium (V) oxide i)

Cr2(SO4)3 j)

strontium phosphide k)

S2F5 l)

manganese (II) phosphate 61) A/n covalent bond forms between two ______; a/n ionic bond forms between a ______and a ______. 62) Circle the pair most likely to form an ionic bond: C-P, O-Cl, K-Cl, S-O 63) Circle the pair most likely to form a covalent bond: Mg-Br, S-O, Li-P, Cr-O

64) What type of bond do the following have in common: AlF3, CaO, LiCl, CuSO4? [Ionic/ covalent] bond.

65) What type of bond do the following have in common: CI4, H2O, N2S2, P2Br4? [Ionic/ covalent] bond. 66) How do atoms form a bond? ______bonds-2 nuclei mutually attract and share electrons. ______bonds transfer electrons to produce ions and electrostatic opposites ______one another. 67) All compounds are [positively charged/negatively charged/neutral]. 68) Atoms bond to obtain a stable ______, full (8 e-s) outer energy level. 69) The ions Br- and Sr+2 both formed a ______octet like __’s outer electron configuration. 70) When an atom has a full ______it is stable like the noble gas. 71) Octet rule-tendency for atoms to lose, gain, or share e-s to have a full outer shell of ___ valence e-s. *Exceptions: H can bond to have only _ electrons in its outer energy level; B can bond to only have _ e-s. 72) Draw Lewis Dot structure, indicate shape and polarity [Polar or NP(nonpolar)] for following:

CI4 BF3 NCl3 H2O F2

tetrahedral/NP 73) Intermolecular forces are ______. 74) As a substance changes state, the temperature [ increases/stays the same/decrease ], b/c added energy is used to overcome intermolecular forces as opposed to increasing ______. Name:______75) As temperature increases, average kinetic energy [ increases/stays the same/decrease] . 76) The greater the attraction between molecules the [ lower/higher ] the melting and boiling points. 77) ______-______forces- interactions b/t polar molecules. Positive end of 1 molecule is attracted to negative end of another molecule 78) ______- special dipole-dipole force between small H bonded to a highly electronegative atom. 79) ______forces are weak dipole forces resulting from temporary uneven e- distribution. 80) London Dispersion (also van der Waals) forces are [ stronger/weaker ] between more massive

molecules. This is why Br2 is a [liquid/solid] at room temperature while I2 is a [liquid/solid] at room T. 81) Number#(1-5) the following in order of decreasing strength (strongest to weakest): __ Covalent bond; __Dipole-dipole; __ Hydrogen bonding; __ Ionic; __ London Dispersion; 82) Which intermolecular force is responsible for water’s unique properties? ______83) The 7 diatomic elements are ___, ____, ____, ____, ____, ____, and ____. 84) When a hydrocarbon produces ______and ______, it is classified as a combustion reaction. 85) The equation A + X  AX is the general equation for a ______reaction. 86) AX  A + X is the general equation for a ______reaction. 87) A + BC B + AC is the general equation for a ______reaction. 88) AB + CD  AD + CB is the general equation for a ______reaction.

89) __Ba(ClO3)2 → __BaCl2 + __O2 ______

90) __C8H18 +__O2 → __CO2 + __H2O ______

91) __Al + __CuSO4 → __Cu + __Al2(SO4)3 ______

92) __Mg(NO3)2 + __Na2SO4 → __MgSO4 + __NaNO3 ______

93) __K + __Cl2 → __KCl ______Gases 94) 6 gas properties: ______. 95) According to ______, gases are composed of small particles far apart relative to size; not attracted to each other; in constant, rapid, random motion; and exert pressure when they collide with container walls. There is no ______lost in “elastic” collisions. 96) The relationship between volume and temperature is ______, as T increases, V ______. 97) The relationship between volume and pressure is ______, as V increases, P ______. 98) The relationship between pressure and temperature is ______, as T increases, P ______. 99) STP stands for ______and ______. Standard temperature is ____; Standard pressure is ______. The standard molar volume of a gas at STP is ______. Reaction Rates Name:______100) Reaction Rate – ______.

101) As temperature increases, reaction rate ______; KClO3 decomposes ______@0°C than 100°C. 102) As concentration increases, reaction rate______; 1M HCl reacts ______than 2M HCl. 103) As surface area increases, reaction rate ______; iron filings rust ______than solid iron nails.

104) Catalysts [raise/lower] activiation energy, Ea,[increasing/decreasing] the rate of

reactions; hydrogen peroxide, H2O2, decomposes ______with KI, a catalyst, than without KI. Equilibrium 105)When forward reactions occur at the same rate as reverse rxns it is ______. 106) ______principle states when a system at equilibrium is disturbed, the system shifts in a way to reduce the stress. Acids and Bases 107) Acids taste _____; bases taste _____. Acids turn litmus paper ____; bases turn litmus paper ____. 108) Arrhenius acids increase______concentrations; Arrhenius bases increase ______concentrations. 109) Bronsted-Lowery acids ______protons, H+; Bronsted-Lowery bases ______protons, H+. 110) pH- ______. 111)The pH of a neutral substance is____. The pH of an acid is______. The pH of a base is ______. 112) An amphoteric species is one that reacts as an______and ______at 25°C? 113) Acids react with bases to produce ______and ______in a neutralization reaction. Thermodynamics 114) Endothermic reactions ______energy; exothermic reactions ______energy. 115) In potential energy (PE) diagrams, endothermic reactions show reactants with _____ PE (more stable) than products; exothermic reactions show reactants with _____PE (less stable) than products. 116)

117) This reaction is an ______rxn.

119) This reaction is an ______rxn. Name:______120) Nuclear 121) ______are involved in chemical reactions; the ______is involved in nuclear reactions. 122) Chemical reactions produce [more/less] energy than nuclear reactions. 123) Organic 124) Carbon has _____ bonding sites. 125) Carbon forms ______bonds, ______bonds, and ______bonds and therefore is able to create many different unique structures important for life. 126) Molar Mass and Percent Composition

127) Barium nitrate, Ba(NO3)2,261.35g/mol,52.55%Ba,10.72%N,36.73%O 128) mole↔gram use molar mass; mole↔particle use avogadro’s #; gram↔particle use 2 steps 24 129) How many formula units are in 6.19moles MgSO4?3.73x10

130) How many moles are in 2.34g CO2?0.0532 3 131) What is the mass of 8.03moles FeCl2? 1.02x10 23 132) What is the mass of 3.58x10 molecules CH4?9.54 22 133) How many formula units are in 28.09g K2SO4?9.707x10

134) What is the empirical formula of a compound with 72.4% iron + 27.6% oxygen? Fe3O4

135) Empirical formula is C2H4S. Molar mass is 179 g/mol. What is the molecular

formula? C6H12S3 136) Stoichiometry: % yield=actual yield/theoretical yieldx100%.

137) Use N2+3H22NH3 for following:

138)How many moles of hydrogen are required to react with 1.327 moles of nitrogen? 3.981 139)How many grams of ammonia are produced form 0.333 moles hydrogen?3.78 140)What mass of hydrogen is needed to produce 55g ammonia? 9.8

141)How many g N2 gas are needed to rxt w 44.8L H2@STP to produce ammonia gas? 18.7

142) Find limiting rxt, theoretical yield, and excess rxt if 14.0g N2 are mixed with 9.0g

H2. N2,17.0,6.0

143)If 14.0g N2 produces 16.1g NH3, what is the percent yield? 94.7

144) Gases: P1V1=P2V2; V1/T1=V2/T2; P1/T1=P2/T2; P1V1/T1=P2V2/T2; PTot=P1+P2+P3…; PV=nRT; R=0.0821(L∙atm)/(mol∙K); R=8.314(L∙kPa)/(mol∙K); 1000mL=1L 145)The volume of a gas is 400.0 mL when pressure is 1.00 atm. At the same temperature, what is the pressure at which the volume of the gas is 2.0 L? *volumes need to be in same unit 0.20 146) A sample of a gas occupies a volume of 752 mL at 25.0°C. What volume will the gas occupy if the temperature increases to 50.0°C, if pressure remains constant? 815 147) A cylinder of helium outside in the sun has a pressure of 5.0atm and a temperature of 50.C. If the cylinder is taken indoors and the pressure decreases to 4.6atm, what’s the temperature? 297 Name:______148)A gas collected @11.0ºC + 710.torr occupies 14.8L. What’s the temperature at 740.torr + 14.7L? 294 149)Calculate the volume of a 0.600mol sample of gas @15.0°C and pressure of 1.10 atm. 12.9

150)What is the mass of CO2 inside a balloon that occupies 11.5L at 30.0°C and 1.79atm? 36.4