Honors Chemistry Midterm Review 2013
Name: __________________________
Date:________________ Period:______
1. Know the proper and safe methods for performing laboratory experiments.
2. Recognize and use the proper definition, instrument and units for the following
measurements:
a. Mass Amount of Matter; balance; g
b. Volume Space taken up; graduated cylinder (ml); LxWxH (cm3)
c. Density Ratio of Mass to volume; Mass/Volume g/ml or g/cm3
d. Time seconds stopwatch
e. Temperature Average Kinetic Energy of molecules; thermometer °C or K
thermometer. Absolute zero is 0K or -273°C K= °C + 273
f. Heat Form of Energy; Joules (J);
g. Amount of substance (mole) 1mole = GFM (gram formula mass)
1 mole = 6.02 x 1023 (atom, mc’s, fu’s)
1mole = 22.4L of a gas at STP
h. Pressure particles colliding on a surface; 1atm = 760mm Hg barometer
i. Area Lx W; ruler cm3
j. Length ruler; cm
3. Metric conversions:
a. ___.01064______________L = 10.64 ml
b. ____7010_____________g = 7.01 kg
c. ____8297.5_____________cm = 82,975 mm
4. Place numbers into proper scientific notation and perform mathematical
operations: use Sig. Figs.
a. ______3.02x108___________ 302,000,000
b. ______7.56x10-9___________ .00000000756
c. ______1.2x1029___________ (1.2 x 10 12) (9.8 x 10 16) =
d. ____1.7x1023_____________ 5.3 x 10 23 / 3.2100 =
5. Identify and use significant figures in calculations:
a. _____3__________ 4.09
b. ______2_________ .0023
c. _____6__________ 7.50000
d. _____.3m2__________ (1.35 m ) (.2 m) =
e. _______850________ 764.3 cm / .90 cm =
f. _______8.5g________ 12.875 g – 4.4 g =
g. ______.047_________
(.00943) (5.0) =
6. Know the properties of elements, compounds, mixtures and solutions. Use this
knowledge to determine (classify) the type of matter.
Element Atoms; simplest form of matter; symbols on p. table
Compound Chemical combination of elements; definite ratio; new properties
Mixtures (2 types) no definite ratio; physical blend; keep original properties
Homogeneous – same throughout (solution)
Heterogeneous- different throughout
7. Compare and contrast mixture, compounds and solutions.
See above. Elements and compounds are both pure and homogeneous. Elements
are one kind of atom; compounds are one kind of “molecule”
8. Identify the location of elements that are metals, nonmetals, metalloids and noble
gases. List their properties.
Metals- Left of the staiarcase except Hydrogen; solids except Hg; good conductors of
heat and electricity/ malleable; ductile; lustrous. Silver/grey in color.
Non-metals – upper right of the staircase; & H; brittle, nonconductors; many gases; sulfur
yellow solid; bromine re-brown liquid; iodine purple; phosphorus is a solid
Metalloids- on either side of staircase; properties on both metals and non-metals;
semiconductors.
Noble Gases- Group 18 (8A); inert non-reactive gases; stable octet ex He (2 valence e-)
9. What are the three states of matter? Describe their physical properties and Kinetic
Molecular properties. Identify the state of matter of each element at STP.
Solid
Attractive forces
packing
motion
v. strong
Moderate to weak
Tightly packed
Medium packed
shape
Vibrate and spin in
place
Definite shape
volume
Definite
compression
Incompressible
State of matter
Liquid
gas
none
Not packed 99%
empty space
Slide past each other Rapid random
straight line motion
Take shape of
Take shape of
container; indefinite container; indefinite
Definite
Take volume of
container; indefinite
Incompressible
Compressible
See periodic table; most element solids; Hg and Br2 are liquids;
STP = standard temperature and pressure; 0°C or 273K; & 1 atm or 760 mmHg of pressure
10. Explain the accomplishments of the following scientists:
a. Democritus; Greek philosopher; idea of the atom; “Atomists”; 2000BC;
not earth, wind fire and water
b. Aristotle; philosopher; vies on physical science; Logic
c. John Dalton- First atomic theory; Atoms solid spheres; Law of Multiple
Proportions
d. Eugene Goldstein- Cathode rays; lead to the discovery of the electrons
1886; discovered protons
e. JJ Thompson- “ Plum Pudding model of the atom
f. Werner Heisenberg- quantum mechanics; uncertainty principle ;
mathematical description of the dual particle like and wave like behavior
g. James Chadwick- 1932; discovery of the neutron
h. Neils Bohr-first step toward quantum mechanical mode; “planetary model”
“electrons traveling in concentric rings around nucleus”
i. Erwin Schrodinger- “electron cloud” quantum mechanical model (modern) “fly
trapped in a jar”
j. Robert Milikan- charge of an electron is exactly -1; Milikan Oil Drop
Experiment
k. Ernest Rutherford; “gold Foil Exp” discovery of the nucleus; positively
charged dense tiny center of the atom. Most of the atom empty space.
“Jimmy Neutron” model
11. Answer the questions below for the element Sodium.
_________Na______________a. Chemical symbol
_________metal______________b. Metal, nonmetal, metalloid, noble gas
_________silver______________c. . Color (silver/grey or other)
________solid__________d. Phase at room temperature
_________yes___________e. Malleable (yes or no)
________yes_______________f. Conductor of heat or electricity (yes or no)
________less_______________g. Is it more or less chemically active than K ?
__________no_____________h. Will it react with iron?
________cation_______________i. Does it form an anion or cation?
________11_______________j. Atomic number
_______22.990 amu________________k. Atomic mass
________23_______________l. Mass Number
_______3________________m. Period
_______1A; 1________________n. Group
________11_______________o. Number of protons
________11_______________p. Number of electrons
________12_______________q. Number of neutrons
________+1_______________r. Charge of the ion it becomes
________Na_______________s. Electron dot structure
________23g_______________t. Mass of one mole.
_______monatomic________________u. Monatomic or diatomic
_______Na+__________v. Formula as an ion
_2.617x1024 atoms______w. Number of atoms in 100.0 grams?
Na2O_____x. Predicted formula of the oxide
______0.9__low______________y. Electronegativity
______1_________________z. How many valence electrons?
__1s22s22p63s1___________aa. Electron configuration
__[Ne]3s1_______________bb. Shorthand configuration
___3s____________________cc. Last sublevel orbital notation
___3,0,0,+1/2_____________dd. Quantum number
____Neon_______________ee. Its ion is isoelectronic to what noble gas?
12. Find the PEN (oh here it is!) for:
Gold Au
Protons
79
Electrons
79
Neutrons
197 – 79 = 118
Carbon-14
6
6
14 – 6 = 8
Strontium ion
Sr+2
38
36
Lost 2 e-
88 – 38 = 50
13. Provide electron dot structures for the following atoms:
14. What do electron dot structures show?
Element; valence electrons; bonding and non-bonding electrons
15. A sample of silver is 52.0% Ag-107 and 48.0% Ag-108, determine the
average atomic mass.
107 x .52 = 55.64
108 x .48 = 51.84
107.48 amu
16. Make the following conversions: 150.00 g of Ar =
a. ____3.7500g Ar___ moles 150g x 1mole
40g
b. ____84.000_______ Liters 150g x 22.4L
40g
c. _____2.2575x1024___________ atoms 150g x 1mole x 6.02x1023 atoms
40g
1mole
17. Make the following conversions: .37 moles of sulfur trioxide = SO3; 32+(3x16) = 80g
a. _____29.6 = 3.0x101_ grams .37moles x 80g
1mole
b. ___8.3__________ Liters
.37moles x 22.4L
1mole
c. ____2.2x1023_____ molecules .37moles x 6.02x1023mc’s
1 mole
d____8.9x1023______ atoms .37moles x 6.02x1023mc’s x 4 atoms
1 mole
1mc
18. How was electromagnetic radiation the key to unlocking the structure of the
atom?
Has properties of both waves and particles; speed = wavelength x frequency
Electron location and movement in atom; Modern Atomic Theory; Quantum
Mechanics
19. Describe how atoms emit electromagnetic radiation.
Electrons jump to higher energy levels. When they drop back down to ground
state they release bundles of energy “photons” in the form of light. Different
amounts of energy = different wavelengths of light ROYGBIV
Emission Spectra can be used to identify elements.
;
h = Planck’s Constant = 6.626 x 10-34 J·s
λ = wavelength measured in meters (m)
c = Speed of Light = 2.998 x 108 m/s
ν = frequency measured in Hertz (1/s or s-1)
E = energy measured in Joules (J)
Use the information above to solve # 20 - 22
20. What is the frequency of a photon of light that has a wavelength of 6.79 x 10-7 m?
λ= 6.79x10-7m c/λ = ⱱ
2.998 x 108 m/s ⱱ = 4.42x1014s
6.79x10-7m
21. What is the energy of a photon of light that has a wavelength of 428 nm?
428nm = 4.28x10-7m
E= hc/λ
6.626 x 10-34 J•s(2.998 x 108 m/s)
4.28x10-7m
E = 4.64x10-19 J
22. What is the wavelength of a photon of light that has an energy of 3.115 x 10-19 kJ?
E= 3.114x10-19kJ = 3.115x10-16J;
λ = hc/E
= 6.626 x 10-34 J•s(2.998 x 108 m/s)
3.115x10-16 J
λ = 6.372x10-10 m
23. Define/determine how the following items are important?
a. E = hf
b. C = wavelength x frequency
c. Hund’s Law- every orbital
electron before they pair up.
in a subshell (s,p,d,f) is occupied by one
d. Hiesenburg Uncertainty Principle- velocity and position of an electron can
not be determined
e. Pauli Exclusion Principle- maximum of 2e- may occupy an orbital and
they must have opposite spin. No 2 electrons can have the same set of
quantum #’s
f. N + l Rule- n=energy level; n=sublevel; n2=orbitals 2n2= #of electrons
Electrons occupy orbital with lower Principal & orbital quantum # totals first.
21.
22.
Name and describe the four Quantum Numbers.
Principal
n
Energy level
Orbital
l
Sublevel
Magnetic
Spin
ml
ms
n= 1,2,3,4……..
l = 0,1,2,3
s,p,d,f
Orientation in space Example for p sublevel; ml = -1; 0; +1
Direction of spin
= +1/2
= - 1/2
23.
How many energy levels do atoms have in the ground state?
There are an unlimited number of energy levels they just may not be
occupied by electrons at ground state.
How many and what type of sublevels do energy levels 1 – 4 have?
24.
Describe the s, p, d, f sublevels in the following terms:
a.
Energy: sublevels, s is the lowest; f is highest; energy increases as energy level #
increases.
b.
Shape: s = spherical; p = 2 lobed; d = 4 lobed; f = 5 lobed
c.
Number of orbitals:
d.
Maximum number of electrons:
25.
How many electrons may occupy a single orbital? Under what conditions?
26.
A neutral atom has the following configuration: 1s22s22p63s23p4
a.
Total number of electrons: 16
b.
Total number of energy levels: 3
c.
Total number of orbitals: n2 = 9
d.
Number of incomplete orbitals: 2
e.
Element: sulfur
27. Find the following for each element listed: electron configuration, shorthand configuration,
last subshell orbital notation, electron dot structure, and quantum numbers.
a. B
1s22s22px1;
[He] 2s22px1;
b. Na
1s2 2s2 2p6 3s1; [Ne] 3s1;
c. Ne
1s2 2s2 2px2y2z2;
[Ne]; 2p
d. O
1s2 2s2 2px2y1z1;
[He] 2s2 2px2y1z1;
e. Ca
1s2 2s2 2p6 3s2 3p6 4s2; [Ar] 4s2;
f. Fe
1s2 2s2 2p6 3s2 3p6 4s2 3d6; [Ar]4s23d6;
2px
; 2,1,-1,+1/2
3s
x
; 3,0,0,+1/2
y
4s
; 2,1,+1,-1/2
z
2p
x
y
z
; 2,1,-1,-1/2
; 4,0,0,-1/2
3d
; 3,2,-2,-1/2
g. W
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d4;
[Xe]6s24f145d4
5d
5,2,+1,+1/2
28. State Mendeleev’s Periodic Law:
Elements arranged by atomic mass show a periodic pattern to their properties.
29. State the Modern Periodic Law:
Elements arranged by atomic number (#p+), show a periodic pattern to their
properties
30. What contributions did the following scientists make?
a. Mendeleev:- 1869; first periodic table arranged by mass
b. Mosely:- x-ray spectra; 1913; modern periodic table; arranged by atomic #
c. Dobereiner: - 1817 Law of triads, Li
Cl
Na
Br
K
I
d. Berzelius:- Atomic weights 1828; chemical formula notation; first to use
symbols and numbers
e. Newlands: -law of octaves 1864; He arranged all the elements known at
the time into a table in order of relative atomic mass. When he did this, he
found that each element was similar to the element eight places further on.
For example, starting at Li, Be is the second element, B is the third and Na
is the eighth element.
31. Using a periodic table locate and give properties for the following:
Metals
Non-metals
Metalloids
Alkali Metals
Alkaline Earth Metals
Halogens
Noble Gases
Transition elements
Lanthanide Series
Actinide Series
Rare Earth Elements
Group A Elements
Group B Elements
Groups 1 – 18
Periods 1 – 7
s block
p block
d block
f block
32. What ion (charge and number) does each group tend to form?
Group A metals the group # is the charge +ions cations
Group B metals; charge may vary but many will be +2
Group A non metals 8- group # = charge
-ions
anion
33. Define terms and list below the periodic trends for the following:
a. Nuclear Charge/Protons:
i.
L to R:
ii.
T to B:
iii.
Element with greatest: N/A
b. Molar Mass/Atomic Mass:-ave. mass of all the isotopes of an element
iv.
L to R:
v.
T to B:
vi.
Element with greatest: N/A
c. Atomic Radii: size; distance from nuclei to nuclei of diatomic element
vii.
L to R:
because energy level same but nuclear charge increasing pulling e- closer
viii.
T to B:
adding another energy level
ix.
Element with greatest: Fr
d. Ionization Energy: energy required to remove an electron
x.
L to R:
because takes more energy to remove e- from smaller atoms; & nonmetals
xi.
T to B:
bigger the atom easier to lose electrons
xii.
Element with greatest: He
e. Electronegativity: attraction for the electrons in a bond
xiii.
L to R:
metals to non-metals who want electrons
xiv.
T to B:
getting bigger and want electrons less
xv.
Element with greatest: F
f. Electron Affinity: Change in energy when an electron is added
xvi.
L to R:
xvii.
T to B:
xviii.
Element with greatest: F
g. Metal Characteristics/Activity: ability to lose electrons
xix.
L to R:
metals to nonmetals who want to gain e- and size decreases
xx.
T to B:
increase in size with each energy level makes it easier to lose e-
xxi.
Element with greatest: Fr
h. Nonmetal Characteristics/Activity: ability to gain electrons
xxii.
xxiii.
metals to nonmetals who want to gain; size getting smaller
greater attraction for eT to B: as size gets bigger its harder to attract electrons
xxiv.
Element with greatest: F
xxv.
xxvi.
xxvii.
xxviii.
L to R:
i. Ionic Radii: metal ions are smaller than their atom;
Non-metal ions are bigger than their atom.
L to R: decrease til 5A then big increase (m to nm) then continue trend of
decreasing as you move across.
T to B: with each energy level added
Element with greatest: Arsenic
j. Shielding Effect: each energy level acts as a shield (block) the nuclear
attraction of the nucleus
L to R: unaffected because no additional energy levels
xxix.
T to B:
increases with each energy level added
xxx.
Element with greatest: period 7
34. Draw the electron dot structure, ionic formula, and empirical formula for a
compound formed from lithium and oxygen.
Electron dot structure;
Empirical formula; Li2O
Ionic Formula: [Li+2 O2-]
35. Draw Lewis Dot structures for the following and determine Bond Type, Molecule
Type; # if sigma and pi bonds, Elements of Symmetry; Hybrid Type; Bond
Angle and Geometry. Note if the molecule demonstrates resonance and/or has
ccb (coordinate covalent bonding)
a. PCl3
Bond Type; P = 2.1; Cl= 3.0 (.9 polar)
EOS; 4
Polar Molecule
dsp3 {pyramidal} bond angle 107° 3 σ bonds
b. SiS2
Bond Type; Si =1.8; S = 2.5 (.7 polar)
EOS; 2
Non polar molecule
sp; linear ; bond angle 180° 2σ bonds & 2π bonds
c. CO2
Bond Type: C = 2.5; O = 3.5 (1.0 polar)
EOS; 2
Non polar molecule
sp; linear bond angle 180° 2σ bonds & 2π bonds
CCl4
d. CCl4
Bond Type: C = 2.5; Cl = 3.0 (.5 polar)
EOS; 4
Non polar molecule
sp3; tetrahedral; bond angle 109.5° 4σ bonds
e. TeO2
ccb
Bond Type; Te = 2.1; O = 3.5 (1.4 polar)
Resonance
EOS; 3 ; polar molecule
sp2; trigonal planar; bond angle 120°
2σ bonds & 1π bond
36. Draw Lewis Dot structures for the following: Note if the molecule demonstrates
resonance and/or has ccb (coordinate covalent bonding). Make sure to circle added ea. PO4-3
Tetrahedral; sp3; 4σ bonds
ccb
b. CO3-2
Trigonal planar; sp2; 3σ bonds & 1π bond
Resonance
37. Write the correct chemical formula for each of the following:
a. ___N2Br6________ Dinitrogen hexabromide
b. __(NH4)2CO3____ Ammonium carbonate
c. __HNO3_________ Nitric acid
d. __Fe(MnO4)3_____ Iron III permanganate
e. __BaF2__________ Barium fluoride
f. __HF___________ hydrofluoric acid
g. __H2SO3________ Sulfurous acid
38. Name the following:
a. ___calcium oxide______ CaO
b. triphosphorous pentoxide_ P3O5
c. _phosphonium sulfite___ (PH4)2 SO3
d. __phosphoric acid_____ H3PO4
e. _Mercuric fluoride (mercury II fluoride)__ HgF2
f. ___hydrosulfic acid_______ H2S
g. ___nitrous acid________ HNO2
39. Vocabulary:
a. Physical change --- Chemical change
b. Homogeneous --- Heterogeneous
c. Element --- Compound
d. Atom --- Molecule
e. Ion --- Isotope
f. Reactant --- Yields --- Product
g. Valence electrons
h. Polyatomic ion
i. Coefficient --- subscript
j. Amu
k. Mole
l. gfm
m. Cation --- anion
n. Catalyst
o. Group --- Period
p. Melting --- Freezing ---Boiling ---Condensation --- Sublimation
q. Solution --- Solvent --- Solute
r. Atomic number --- Atomic mass --- Mass number ---gfm
s. Malleable --- Ductile
t. Diatomic molecules