Download MYP Chemistry: Final Review

MYP Chemistry: Final Review
Unit 1: Conservation of matter
 Polyatomic ions (formula and charge)
 Writing formulas
 Naming compounds
 Reading and writing chemical equations
 Balancing equations
 Br i n cl h o f
 Classifying reactions (synthesis, decomposition,
combustion, single replacement, double replacement)
 Predicting products of a reaction
Unit 2: The Atom
 Basic atomic structure; locations/charges/masses of
subatomic particles
 Atomic symbols, calculating atomic mass
 Wave relationships: wavelength, frequency, and energy
 EM spectrum, continuous spectrum vs. Bright line
spectrum
 Bohr model, relationship to bright line spectra
 Electron dancing, identifying elements with spectra,
 Electron configuration
 Organization of the Periodic Table – Periods and Groups

Names of common families: alkali metals, alkaline earth
metals, boron family, carbon family, nitrogen family,
oxygen family, halogens, noble gases
Unit 3: Bonding
 Periodic trends: atomic radius, ionization energy,
electronegativity, reactivity
 Ionic bonding: electron transfer
 Covalent bonding: electron sharing,
 Lewis dot structures, exceptions to octet rule,
double/triple bonds, ions
 Shapes of molecules (VSEPR), bond angles, resonance,
effect of shape on molecular polarity
 Intermolecular Forces
Unit 4: Properties of Matter
 Matter definition
 The mole (moles/molecules)
 Atomic mass; formula mass (g/mole)
 Solutions
 Molarity Calculations
PRACTICE PROBLEMS
Nomenclature:
Identify each compound as either ionic or molecular. Then name it and determine formula mass.
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Compound
MgO
SO4
Fe2S3
N 2O 5
BaSO3
Zn(SO4)
CCl4
SF6
Ni3(PO4)2
Ionic/Covalent
I
C
I
C
I
I
C
C
I
Name
magnesium oxide
sulfur tetroxide
iron (III) sulfide
dinitrogen pentoxide
barium sulfite
zinc (II) sulfate
carbon tetrachloride
sulfur hexafluoride
nickel (II) phosphate
Molar Mass
40.3
96
208
108
217.3
161.4
154
146
366.1
Identify each compound as either ionic or molecular. Then write a correct formula and determine formula mass.
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Compound
magnesium sulfate
nitrogen triiodide
lead (II) phosphate
ammonium nitrate
dichlorine monoxide
carbon dioxide
aluminum chloride
chromium (III) oxide
potassium iodide
diphosphorus tetraoxide
Ionic/Covalent
I
C
I
I
C
C
I
I
I
C
Name
MgSO4
NI3
Pb3(PO4)2
NH4NO3
Cl2O
CO2
AlCl3
Cr2O3
KI
P2O4
Molar Mass
120.3
395
811
80
87
44
133.5
151.99
166
126
Reactions and equations:
For each of the following: write a balanced equation and classify the reaction.
 sodium reacts with aqueous aluminum carbonate 6Na + Al2(CO3)3  3Na2CO3 + 2Al (SR)
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aqueous sodium sulfide reacts with aqueous barium nitrate Na2S + Ba(NO3)2  BaS + 2NaNO3(DR)
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copper (II) carbonate reacts to form copper (II) oxide and a gas CuCO3 CuO + CO2 (DEC)
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potassium reacts with iodine 2K + I2  2KI (SYN)
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ethane (C2H6) burns in air 2C2H6 + 7O2  4CO2 + 6H2O (COMB)
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barium reacts with zinc (II) chloride Ba + ZnCl2  BaCl2 + Zn (SR)

sodium carbonate reacts with zinc (II) fluoride Na2CO3 + ZnF2  2NaF + ZnCO3 (DR)
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lead (II) oxide decomposes 2PbO  2Pb + O2 (DEC)
Atomic Structure:
 Sketch a diagram of an atom including locations, charges, and masses of all three subatomic particles.
electrons (-) 1/1840 amu
Nucleus:
- protons (+) 1 amu
- neutrons (0) 1 amu

Define (or write the equation that defines them): atomic number, mass number, average atomic mass, atom, ion, and isotope.
Atomic Number = # protons = Element;
Mass # = #neutron + #protons;
ion: a charged atom;
isotope: similar elements (same number of protons) with a different mass # (different # of neutrons);
average atomic mass: weighted average of all the isotopes of an atom;
atom: basic unit of a chemical element

Provide an example of an ion. Cl-1 or Na+1
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Provide an example of two or more isotopes:
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What two subatomic particles contribute to an atom’s mass? proton and neutron
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What two subatomic particles contribute to an atom’s charge? proton and electron
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Fill in the chart:
Isotope
14
C and 12C
# of
protons
# of
neutrons
# of
electrons
26
31
26
6
8
6
6
6
6
17
18
18

Calculate the atomic mass of a sample that contains 20% 6 Li and 80% 7Li. (0.20 * 6) + (0.80 * 7) = 6.8 amu
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What experiment did J.J. Thomson run? What did he discover? Cathode Ray Tube experiment. He discovered electrons and
that they were negative
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What experiment did Ernes Rutherford run? What did he discover? Ernest Rutherford ran the gold foil experiment. He
discovered the nucleus and that atoms are mostly empty space
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How did Bohr contribute to Rutherford’s atomic model? Bohr added energy levels around the nucleus (like planets orbiting
the sun)
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Determine the average atomic mass of Gadolimium if the following information is known:
Isotope
Percent
abundance
Gadolimium – 158
157.84
25.34
Gadolimium - 160
159.84
50.00
Gadolimium – 162
161.84
24.66
(158*.2534) + (160 * .5000) + (162*.2466) = 157.00 amu
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Mass (amu)
What is the electron configuration of sodium? Of chlorine?
Na: 1s22s22p63s1
Cl: 1s22s22p63s23p5
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What element has the electron configuration 1s22s22p3? … [Ne]3s23p3?
N
P
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How are elements in the same group (column) related? How are the alkali metals all related? The noble gases?
All have the same final electron configuration number; all have same number of valence electrons
Alkali Metals: end in s1 configuration, have 1 valence electron
Noble gases: end in s2p6, have 8 valence electrons
Nature of Light:
 As wavelength gets shorter, frequency increases
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As wavelength gets shorter, energy increases
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What is the electromagnetic (EM) spectrum? What is the highest energy wave? Lowest? The range of all known light, the
lowest being radio waves and the highest being gamma rays. Visible light (ROYGBV) is a tiny portion of the spectrum in
the middle
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What is the difference between a bright line spectrum and a continuous spectrum? Continuous spectrum contains all
wavelengths (ROYGBV) like a rainbow. Bright line spectrum shows discrete wavelengths like red or blue or green, but not
all the colors
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Explain how electron movement between energy levels produces photons of light. When an atom is excited by an energy
source, electrons jump to higher energy levels. Since they are (-) charged and are attracted to the (+) nucleus, they jump
back down to their original level. This jump down releases energy in the form of a photon of light

Consider this diagram of an atom with arrows representing electron movement.
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a) Which two arrows correspond to energy absorption by the atom? B and D
b) Which two arrows correspond to energy emission by the atom? A and C
c) If violet and green light are produced by the movement illustrated here, which arrow represents emission
of violet light? A green light? C
How do waves of red light and blue light differ with respect to frequency? wavelength? energy?
Blue – higher energy/frequency; shorter wavelength
Red – lower energy/frequency; longer wavelength
Periodic Trends:
 Define the terms electronegativity, ionization energy, atomic radius. Describe their trends on the periodic table.
Electronegativity: tendency for and atom to gain an electron;
ionization energy: energy required to remove an electron from an atom; atomic radius: size of atom;
Electronegativity and Ionization Energy decrease down a group and increases across a period.
Atomic radius increases down a group and decreases across a period
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Sr
is the 5th period alkaline earth metal
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Cl
is the 3rd period halogen
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Br
is the most reactive nonmetal in the 4th period
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K
has the largest radius of the 4th period
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Ra
is the largest atom with two valence electrons
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Be
is the least reactive alkaline earth metal
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Cl
has 18 electrons and 17 protons
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Mg
has 2 valence electrons in the 3rd energy level
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Be
has the greatest ionization energy of the alkaline earth metals
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S
is the element in oxygen family with 3 energy levels
-1
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Na
is the 3rd period element with lowest ionization energy
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Ne
is the noble gas with a 2nd energy level valence shell
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Rb
is the 5th period atom most likely to lose an electron
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Xe
is the 5th period atom least likely to react
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4d
For elements 39 through 48, additional electrons are added to the ______ sublevel
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Se
Which element has the greatest radius: S Cl Se Br
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Al
is the third period element which is most likely to form an ion with a +3 charge
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At
The halogen with the highest melting point
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nonmetals
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metals
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Rb
Write the symbol that has the largest radius Rb Sr K Ca
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Rb
Write the symbol that has the greatest reactivity Rb Sr K Ca
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Cl
Write the symbol that is the most reactive Se Br S Cl
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Cl
Which has the highest electronegativity? Se Br S Cl
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Rb
Which has the lowest electronegativity? Rb Sr K Ca
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Why do elements form chemical bonds? to increase their stability by completing their octet
react by gaining electrons.
react by losing electrons.
Covalent Bonding:
For each of the following: a) draw a dot structure b) give the shape of the molecule c) does the mode have resonance?
 H2O
 BCl3
- trigonal planar
- nonpolar
- 120°
- bent
- polar
- 104°
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- linear
- nonpolar
- 180°
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C2H4
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HCl
- linear
- polar
- 180°
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- trigonal biplanar
- nonpolar
- 120° and 180°
NH3
- trigonal pyramidal
- polar
- 107°
NH4+
+ - tetrahedral
- nonpolar
- 109.5°
CO2
CH2O
- trigonal planar
- polar
- 120°
CH4
- tetrahedral
- nonpolar
- 109.5°
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F2
- linear
- nonpolar
- 180°
C2H2
- linear
- nonpolar
- 180°
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What are some distinctive properties of ionic compounds?
Brittle, high melting/boiling points, very stable, conducts electricity when dissolved, solid at room temperature
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What makes an ionic bond different than a covalent bond?
Ionic bonds created by a complete transfer of electrons, electrostatic attraction between opposite ions
Covalent bond created by sharing electrons, evenly (covalent) or unevenly (polar covalent)
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What is VSEPR theory?
Valence Shell Electron Pair Repulsion: bond angles increase with the increasing number of valence electron pairs
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Draw a Lewis dot structure for each of the following molecules: O 2, H2O, CH2Cl2, NI3
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In the question above, which molecules are polar? Which are nonpolar? Assign geometries to each.
O2 – nonpolar, linear
H2O – polar, bent 105
CH2Cl2 – polar, tetrahedral
NI3 – polar, trigonal pyramidal
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What are the three types of intermolecular forces? What type(s) of molecules is each one present in?
London, van der Waals, Dispersion – nonpolar molecules, attracted by electrons
Dipole – polar molecules, attracted due to slightly negative charge on one molecule to slightly positive charge on
another molecule
Hydrogen Bond – very polar molecules, H bonded to N,O,F creates a H-bond to another very polar molecule
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What are some properties of the types of compounds?
Covalent – s, l, g at room temperature, low to high melting points, low to high solubility, poor to non-conducting
Network – brittle, high melting/boiling point
Metals – good conductors, ductile, malleable, solid at room temperature (except Hg)
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Describe the differences (on the molecular level) between solid, liquid, and gas phases.
Solid – molecules vibrating in set space, low kinetic energy
Liquid – molecules vibrate while flowing within container, medium kinetic energy
Gas – molecules have highest kinetic energy, move rapidly and spastic
Matter:
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Define the following terms: matter, pure substance, homogeneous mixture, heterogeneous mixture, element, atom, and
compound.
Matter: anything that has mass and takes up space;
pure substance: matter that is uniform and has a definite composition;
homogeneous mixture: two or more substances physically blended and is completely uniform in composition;
heterogeneous mixture: two or more substances physically blended and is not uniform in composition;
element: simple, pure substance, atoms of the same type;
compound: two or more elements chemically bonded.
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What are the two types of mixtures? Describe each and say how we can tell them apart.
Homogenous and heterogeneous.
Homogeneous is consistent throughout like Kool-Aid dissolved in water.
Heterogeneous is not consistent throughout like salsa.
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Classify each of the following as an element, compound, heterogeneous mixture, or homogeneous mixture.
a. Gold (Au) element b. Kool-Aid completely dissolved in water homogeneous mixture c. Sodium bicarbonate
(H2CO3) compound d. Sulfur (S) element e. NaCl compound f. Salsa heterogeneous mixture
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a.
d.
a.
Classify each of the following changes as either chemical or physical.
Cutting a sheet of aluminum foil into 4 pieces Physical b. Burning of coal Chemical c. Cooling a liquid until it freezes
Physical
A white solid and sulfuric acid are mixed and an orange gas is produced Chemical e. Dissolving sugar in tea Physical
A piece of copper is hammered into a thin sheet Physical
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Give three examples of physical properties and three examples of chemical properties.
Physical Properties: density, melting point, boiling point, smell, texture, color, etc. Chemical Properties:
combustibility, acid/base, flammability, oxidation
Mole calculations
5.6𝐿∗1 𝑚𝑜𝑙𝑒
0.25 𝑚𝑜𝑙𝑒𝑠∗37.996𝑔
 Find the mass in grams of 5.6L molecules of F2.
= 0.25 𝑚𝑜𝑙𝑒𝑠
= 9.5 𝑔
22.4𝐿
1 𝑚𝑜𝑙𝑒
𝟒𝟖𝟏𝒈𝑨𝒓∗𝟏 𝒎𝒐𝒍𝒆
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Find the number of moles of argon in 481 g of argon.
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Calculate the number of moles in 6.0 grams of carbon.
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Penicillin F has the formula C14H20N2SO4. How many moles of this medicine are in a 2.0 g dose?
𝟑𝟗.𝟗𝟒𝟖𝒈𝑨𝒓
𝟔.𝟎𝒈𝑪∗𝟏 𝒎𝒐𝒍𝒆
𝟏𝟐.𝟎𝟏𝟏𝒈𝑪
= 𝟏𝟐. 𝟎𝟒 𝒎𝒐𝒍𝒆 𝑨𝒓
=. 𝟓𝟎 𝒎𝒐𝒍𝒆 𝑪
𝟐.𝟎𝒈𝑷𝒇∗𝟏𝒎𝒐𝒍𝒆
𝟑𝟏𝟐.𝟑𝟑𝒈
=
𝟎. 𝟎𝟎𝟔 𝒎𝒐𝒍𝒆𝒔
𝟏𝟎𝟎.𝟎𝒈𝑶𝟐∗𝟏𝒎𝒐𝒍𝒆

Calculate the volume of 100.0 g of oxygen gas.
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How many moles of sodium are in 1 mol of Na2O? 2 mol Na
𝟑𝟏.𝟗𝟗𝟖𝒈𝑶𝟐
= 𝟑. 𝟏𝟑 𝒎𝒐𝒍𝒆 𝑶𝟐
𝟑.𝟏𝟑𝒎𝒐𝒍𝒆𝑶𝟐∗𝟐𝟐.𝟒𝑳
𝟏𝒎𝒐𝒍𝒆
= 𝟕𝟎. 𝟎𝟎 𝑳 𝑶𝟐
Solutions
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Define the terms solute, solvent and solution.
Solute: what’s being dissolved
Solvent: the dissolver
Solution: mixture
Detail the differences between a heterogeneous mixture and a homogeneous mixture. Which category does a solution fall
into?
Heterogeneous mixture has visible particles that identify the different substances in the mixture whereas the particles
in a homogeneous mixture are not distinguish and will not settle out of solution.
Homogeneous solution – no visible particles
Heterogeneous collide – small particles that stay suspended
Heterogeneous suspension – visible particles that settle out of solution easily

Define the term concentration. List a few of the units and formulas used to determine concentration (particularly
molarity).
Concentration: the amount of solute in a solution
𝒎𝒐𝒍𝒔
𝑴=
𝑳𝒊𝒕𝒆𝒓𝒔

What is the molarity of a solution containing 2.5 grams of HCl in 3.5 liters of solution?
𝟐.𝟓𝒈𝑯𝑪𝒍∗𝟏 𝒎𝒐𝒍𝒆
.𝟎𝟔𝟗 𝒎𝒐𝒍𝒆𝒔
= . 𝟎𝟔𝟗 𝒎𝒐𝒍𝒆𝒔 𝑯𝑪𝒍
= 𝟎. 𝟎𝟐𝑴
𝟑𝟔.𝟒𝟔𝒈𝑯𝑪𝒍

𝟑.𝟓𝑳
How many moles are contained in 3.4 liters of a 2.5 M solution?
𝒎𝒐𝒍𝒆𝒔
(𝟐. 𝟓 ∗ 𝟑. 𝟒) = 𝟖. 𝟓𝒎𝒐𝒍𝒆𝒔
𝟐. 𝟓𝑴 =
𝟑.𝟒𝑳