1. A) Heterogeneous B) Heterogeneous C) Homogeneous D) Heterogeneous E) Homogeneous

Exam Review Answers

1. a) heterogeneous b) heterogeneous c) homogeneous d) heterogeneous e) homogeneous

2. a) chemical b) physical c) chemical d) physical e) chemical f) physical g) physical

Element / Symbol / Atomic # / Mass # / # Neutrons
nitrogen-15 / / 7 / 15 / 8
neon-22 / / 10 / 22 / 12
beryllium-9 / / 4 / 9 / 5

4. (a) (b) (c) (d)

5. (a) selenium (b) strontium (c) iodine(d) sulfur(e) astatine(f) sodium(g) radon

6. (a) rep. (b) rep. (c) trans. (d) inner trans. (e) trans. (f) metalloid (g) rep. (h) trans. (i) trans. (j) rep

7. a) H+ b) Mg2+ c) S2- d) I 1- e) Al 3+

8. a) Valence electrons – electrons within the valence shell (highest numbered energy level.

b) anion – a particle that has a negative charge do to a greater number of electrons than protons.

c) cation – a particle that has a positive charge do to a greater number of protons than electrons.

d) ionization energy – the measure of energy required to remove an electron from an atom in the gaseous state.

e) electronegativity – a measure of an atom’s ability to attract bonding electrons.

f) octet rule – atoms become stable by attaining a complete valence shell.

g) ionic bond – atoms gain or lose valence electrons to attain a full valence shell and the resulting oppositely charged ions attract each other.

h) covalent bond – atoms share valence electrons to attain a full valence shell and the resulting mutual attraction to the bonding electrons holds the atoms together.

i) polar bond – covalent bonds that are formed with unequal sharing of bonding electrons shown through an electronegativity difference for the atoms within the range of 0.4 to 1.7

j) nonpolar bond – covalent bonds that are formed with equal sharing of bonding electrons shown through an electronegativity difference for the atoms within the range of 0 to0.4

K 1s22s22p63s23p6 4s1K+ 1s22s22p63s23p6 S 1s22s22p63s23p4 S2- 1s22s22p63s23p6

Ca 1s22s22p63s23p64s2Ca2+ 1s22s22p63s23p6 Ar 1s22s22p63s23p6

10.

(a) Dalton’s Atomic TheoryAll matter is composed of small indivisible particles

Atoms of an element are identical in size shape and mass

Atoms of one element are different than atoms of another element

Chemical reactions are the rearrangement of atoms (number of atoms before and after the same

(b) Thomson’s model of the atomatom small particle made up of positive matter with negative electrons imbedded.

(c) Rutherford’s model of the atomatom viewed as a small massive positively charged nucleus surrounded by extra-nuclear electrons.

(d) Bohr’s model of the atomatom viewed as a small massive positively charged nucleus surrounded by extra-nuclear electrons found in orbit determined by their amount of energy (low energy small orbit; large orbit higher energy).

(e) Quantum Mechanical model atom viewed as a small massive positively charged nucleus surrounded by extra-nuclear electrons found in energy levels that contain sublevels that contain orbitals of definite shapes (shapes are based on the probable location of the electrons).

11.

Element / # Valence e- / Dot Structure / Ion Formed / Anion/Cation / Noble Gas
O / 6 / 6 dots-2 lone pairs, 2 dots / O2- / Anion / Ne
Na / 1 / 1 dot, no lone pairs / Na+ / Cation / Ne
P / 5 / 5 dots-1 lone pair, 3 dots / P3- / Anion / Ar
Ca / 2 / 2 dots, no lone pairs / Ca2+ / Cation / Ar

12.ionic compounds contain metals and nonmetals and are formed when the metal atoms loses valence electron(s) to the nonmetal (atoms end up with a complete octet). The metal cation attracts the nonmetal anion. The ions arranged in a crystal lattice which maximizes the neutralization of the ionic charges.
molecular compounds contain nonmetal with a nonmetal and are formed when the atoms share valence to attain a complete octet. The mutual attraction to these shared (bonding) electrons holds the atoms together to form a molecule – an electrically neutral particle. When the bonding electrons are shared equally a nonpolar bond is formed. When the bonding electrons are not shared equally a polar bond is formed.

13.

a)NaIionicsodium iodide

b)NOmolecularnitrogen monoxide

c)GaBr3ionicgallium bromide

d)S4N2moleculartetrasulphur dinitride

e)ZnCl2ioniczinc chloride

f)SnF4ionictin(IV) fluoride

g)P2O5moleculardiphosphorus pentoxide

14.

a)hydrochloric acidHCl(aq)

b)strontium nitrateSr(NO3)2

c)calcium chlorideCaCl2

d)ammoniaNH3

e)sulphur trioxideSO3

f)lithium sulphateLi2SO4

g)calcium hydroxideCa(OH)2

h)ammonium sulphate(NH4)2SO4

i)lead (II) sulfitePbSO3

j)copper (I) sulphideCu2S

k)aluminum oxideAl2O3

l)magnesium bromideMgBr2

m)hydrogen gasH2(g)

n)nitrogen dioxideNO2

15. (a) N (b) Au (c) K (d) Ga

16. (a) O (b) Ag (c) Ca (d) Ge

17. (a) O (b) Ag (c) Ca (d) Ge

18. (least) Fr, Te, Ge, Mg, C, F (most)

19. (a) C-F(b) P-N(c) I-F(d) C-N

20. Most:FLeast:Fr

21.

H – Cl

δ+ δ–

δ+H – O δ–

δ+H

22.

(a) ionicMg2+ and Cl-

(b) covalentbentpolar

(d) covalentnonpolar

(e) ionicAl3+ and S2-

(f) covalentnonpolar

(g) ionicNa+ and N3-

(h) ionicCa2+ and O2-

(i) covalentpolar

(j) covalent polar

23. Types of intermolecular forces:

Ionic bonding very strong attractions between oppositely charged ions leads to the formation of rigid solids that shatter when struck by a force and have high melting temperature.

Dipole – dipole attraction is a weaker attractive force that occurs between opposite partial charges for molecular compounds containing polar covalent bonds (unequal sharing of bonding electrons).

Van der Waal’s (dispersion) forces are very weak forces that are a result of the uneven distribution of the electrons in the molecular orbital. This separation of charge is momentary (short lived) so the attraction to other momentary charges are very weak. The bigger the molecular the longer the separation of charge exists; the greater the dispersion force of attraction.

24.

1CF4(l) → 1C(s) + 2F2(g)
1H2SO4(aq) + 1KOH(aq) → 1KHSO4(aq) + 1H2O(l)
1ZnCl2(aq) + 1H2(g) → 1Zn(s) + 2HCl(aq)
2SO2(g) + 2H2O(l) + 1O2(g) → 2H2SO4(aq)
2Li(s) + 2H2O(l) → 2LiOH(aq) + 1H2(g)
1H2CO3(aq) → 1H2O(l) + 1CO2(g)
1Na2SO4(aq) + 1BaCl2(aq) → 1BaSO4(s) + 2NaCl(aq)
2CH3OH(l) + 3O2(g) → 2CO2(g) + 4H2O(g)

25. a) decomposition b) double replacement c) single replacement d) composition

e) single replacement f) decomposition g) double replacement h) combustion

26.

a) 1Pb(NO3)2 (aq)+ 2NaI(aq) 1PbI2(s) + 2NaNO3(aq)double replacement

b) 1CaCO3(s)1CaO(s) + 1CO2(g)decomposition

c) 1C3H8 + 5O23CO2 + 4H2Ocombustion

d) 1Cu(s) + 2AgNO3(aq)  1Cu(NO3)2(aq)+ 2Ag(s)single replacement

e) 2Na(s) + 1Cl2(g) 2NaCl(s)composition

27.

A) CO2B) Na2SO4

1 x C = 1 x 12.01g = 12.01g2 x Na = 2 x 22.99g = 45.98g

2 x O = 2 x 16.00g = 32.00g1 x S = 1 x 32.06g = 32.06g

4 x O = 4 x 16.00g = 64.00g

1 mole = 44.01g1 mole = 142.04g

28.

(a) 25.0 g of NaCl x 1 mole = 0.428mole NaCl(b) 125 g of H2SO4 x 1 mole = 1.27mole H2SO4

58.44g98.08g

29.

(a) 1.70 moles of K2O x 94.2g = 160.g(b) 0.25 moles of KCl x 74.548g = 18.64 g

1 mol1 mole

30.

(a) 10.0 L H2(g) at STP x 1 mole = 0.446mole(b) 48.5 L O2(g)) at STP x 1 mole = 2.17 mol

22.4 L22.4 L

31.

(a) 15.0 L Ne(g) at STP x 1 mole = 6.25mole x 20.18 g = 126g

22.4 Lmole

(b) 44.0 L He(g) at STP x 1 mole = 1.96 mole x 4.00g = = 7.86g

22.4 L mole

32.N2(g) + 3H2g)→ 2NH3(g)

a)H2 = 3 = xmolexmol = 6mol H2

N2 1 2mole

50.0gN2 x 1 mole = 1.78 moleH2 = 3 = x x =5.34 mole H2 x 2.02 g = 10.8g H2

28.02gN2 1 1.78 1mole

33.NaOH(aq) + HCl(aq) → NaCl(aq) + H2O(l)

14.0g NaOH x 1 mole= 0.35mol NaOHNaCl = 1 = x x = 0.35mol NaCl x 58.44g = 20.5g

40.00gNaOH 1 0.351mole

34.and.

The coefficient states the number of formula units or molecules present (number of moles).

A subscript shows the number of ions or atoms present in the formula for the compound.

35.3NO2(g) + H2O(l) → 2HNO3(aq) + NO(g)

36.04g H2O x 1 mole = 2.00mole H2O NO2 = 3 = x x = 6moleNO2 x 22.4mole = 134L

18.02gH2O 1 2.00moleat STP