Midterm Review

Midterm Review
Chemistry Level 2
Chapter 1
• What is chemistry?
– The field of science that studies the composition
and structure of matter
• What is matter?
– Anything that has mass and occupies space
• Be able to describe each step of the scientific
method. Provide a couple examples where
you used the scientific method to solve a
problem in your life and label each part.
– State a Problem, Hypothesis, Experiment, Reach a
Chapter 1 - Continued
• Theory v. Scientific Law --- What is the
difference? Think of examples for both and
identify the limitations. Can they be
proved/disproved supported/falsified? Which
one is more powerful than the other?
– Law summarizes observations but theories
provide an explanation
– Limitations:
• A scientific Theory CAN NEVER BE PROVED
Chapter 1 - Continued
• What are variables? What step of the
scientific method involves manipulating
– Variables are factors that can change
– Experimenting
• In large, why has science advanced as much as
it has over the years?
– Advancements in science has taken part largely
because of advancements in technology
Chapter 2
• Name some examples of physical properties.
– Mass, volume, density, boiling point, melting point,
solubility, phase (solid, liquid or gas), color, texture,
hardness, etc…
Chapter 2 (continued)
• Define the particle arrangement of a solid,
liquid and gas. What can you say about their
shapes and the volume they occupy?
– Solids have an orderly arrangement of particles
with a definite shape and a definite volume
– Liquids have a less orderly arrangement of
particles with no definite shape and definite
– Gases have a random arrangement of particles
with no definite shape and no definite volume.
Chapter 2 (continued)
• Identify physical properties from chemical
properties. For example zinc metal is hard,
silver in color, is easily shaped and reacts
vigorously with hydrochloric acid. What are
some physical and chemical properties of
– Physical
• Metal, hard, silver, malleable
– Chemical
• Reacts with hydrochloric acid
Chapter 2 (continued)
• What is the difference between a physical and
chemical change? How do you know when
one occurred? Think of examples.
– A physical change is one in which the substance
still retains its identity
• Evaporating, breaking, bending, cutting, etc…
– A chemical change is one in which its chemical
identify has changed
• Rusting, burning, explosion, decaying, etc…
Chapter 2 (continued)
• What are symbols? What are formulas?
When should you capitalize letters? What are
– Symbols represent elements
– Formulas represent compounds
– The first figure in a properly written chemical
symbol always is capitalized.
Chapter 2 (continued)
• An object has a volume of 12mL and a mass of
3.5g. What is the object’s density?
– 0.29 g/mL
• An object has a mass of 357g and volume of
75mL. What is the object’s density?
– 4.8 g/mL
• An object has a mass of 4.2g and a volume of
19.3mL. What is the object’s density?
– 0.22 g/mL
Chapter 3
• What is an advantage of the metric system compared
to the English system?
– Based on a power of 10
• What is the SI unit of mass? What is the SI unit for
– kilogram – mass
– Kelvin - temperature
• Define weight. Does weight change based on location?
– the pull on a given mass by gravity
– the weight of an object depends on its location
• Which temperature scale has no negative numbers?
– Kelvin
Chapter 3 (continued)
Convert the following:
50oC = ___323_______K
-25oC = ___248___K
Count the number of significant figures in these
50,400 mg = _____3_______
0.00046 kg = _____2_______
923.110 mg = _____6______
1.0045 L = _______5_______
Chapter 3 (continued)
Round problems 1-4 to two significant figures
and write the answer in scientific notation.
50,400 mg = ____5.0 X 104 mg________
0.00046 kg = ____4.6 X 10-4 kg_______
923.110 mg = ____9.2 X 102 mg______
1,045 L = _______1.0 X 103 L______
Chapter 3 (continued)
Review rules regarding adding/subtracting and multiplying/dividing
1. What is the result of adding
2.01g +4.5g = 6.5 g
2. What is result of subtracting
4.356m – 3.6m = 0.8 m
3.What is the result of multiplying 5.2 X 102 by 1.367 X 10-4m = 7.1 X 10-2 m
4.What is the result of dividing
4.3 X102 by 2.0 X 104g = 2.2 X 10-2 g
Chapter 3 (continued)
• Define density.
– Ratio of an object’s mass to its volume
– Density is found by dividing mass by volume
• volume = 90.0 mL
• mass = 55 g
• mass = 7.0 g
density = 0.70g/mL
density = 12.3 g/mL
volume = 9.1 cm3
mass = 63 g
volume = 4.5 mL
density = 0.77 g/cm3
• What is the density of an object that has a
mass of 5.6g and a volume of 25cm3?
– 0.22 g/cm3
Chapter 4
Fill in the table below:
Positive (+)
1 amu
In nucleus
No charge
1 amu
In nucleus
Negative (-) 0 amu
Chapter 4 (continued)
• What does the atomic number of an element
– # of protons
• What does the mass number of an element
– number of protons + neutrons
• If an atom has an atomic number of 36 and a
mass number of 84 how many protons,
neutrons and electrons are present in this
– 36 Protons, 36 electrons, and 48 neutrons
Chapter 4 (continued)
• What does the number 13 represent in carbon-13?
– Mass Number (number of protons and neutrons)
• How do you calculate the number of neutrons in an
Mass Number - # of Protons = # of Neutrons
• Who was Democritus?
– One of the first to suggest the idea of atoms
(460 -370 B.C.)
• What is the unit that is used to measure a weighted
average atomic mass?
– Amu (atomic mass unit)
Chapter 4 (continued)
• What is an isotope?
– Isotopes of the same element have different
numbers of neutrons
• What is an ion?
– Ions are atoms that have number of protons not
equal to electrons
• How many protons, neutrons and electrons
are present in this neutral atom? 3919K
– 19 protons, 19 electrons, and 20 neutrons
Chapter 5
• Describe the quantum mechanical model of
an atom?
– Based on the probability of finding an electron
• Which scientist developed the quantum
mechanical model of an atom?
– Erwin Schrödinger
• What are the shapes of an s and p orbitals?
– s is a spherical shape
– p is a dumbbell shape
Chapter 5 (continued)
• What is a principal energy level, sublevel and
atomic orbital?
– Principal energy level contains one or more types
of sublevels
– Sublevels are atomic orbital shape represented by
the following letter s, p, d and f
– Atomic orbitals are the regions of space where
electrons are located in the different sublevels
S sublevel
P sublevel
Principal Energy Level 1
Principal Energy Level 2
Chapter 5 (continued)
• What is the maximum number in each s, p, d
and f orbitals?
– S (1 orbital), p (3 orbitals), d (5 orbitals), f (7
• What types of atomic orbitals are in the 1st,
2nd and 3rd principal energy levels?
– 1st (s atomic orbitals)
– 2nd (s and p)
– 3rd (s, p, and d)
Chapter 5 (continued)
• If the spin of one electron is clockwise in an
orbital the spin on the second electron must
be counterclockwise?
• Using the Aufbau diagram what orbital would
come after 3p?
– 4s
• What is the number of electrons in the
outermost energy level of sulfur?
Chapter 5 (continued)
• What happens when an electron moves from
a lower to higher energy level?
– absorb a quantum of energy
• When an electron moves from a higher to
lower energy level what happens?
– lose a quantum of energy
• When does emission of light occur?
– Drops from a higher to lower energy level
Chapter 5 (continued)
• Which of the following quantum leaps would
be associated with the greatest energy
– n = 4 to n=2
– n = 5 to n= 1
Chapter 5 (continued)
• How are frequency and wavelength of light
– They are inversely proportional to each other
• Which variable is directly proportional to
– frequency
• Define a photon.
– Quanta of light
Chapter 5 (continued)
• What is the wavelength of a radiation with a
frequency of 5 X 1018 Hz (s-1)? (c = 3.00 X 108
6 X 10-11 m
• What is the frequency of light having a
wavelength of 4.2 X 10-8m?
7.1 X 1015 m
• What is the energy of a photon with a frequency
of 2.5 X 1015 Hz? [E = h v] (h = 6.63 X 10-34 J s)
1.7 X 10-18 J
Chapter 5 (continued)
• What are the electron configurations for the
following elements:
– Na:
– Ar:
– Zn:
– As:
– Br:
– Sr:
1s2 2s2 2p6 3s1
1s2 2s2 2p6 3s2 3p6
1s2 2s2 2p6 3s2 3p6 4s2 3d10
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p3
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2
Chapter 6
• Another name for the representative
elements is Group A elements. Where are
these elements located on the periodic table?
– In the s and p sublevels
• Who was the first scientist to arrange the
elements according to similar chemical and
physical properties in order of increasing
atomic mass?
– Dmitri Mendeleev
Chapter 6 (continued)
• What is characteristic of the electron
configurations of noble gases?
– The highest occupied s and p sublevels are
completely filled
• Which subatomic particle plays the greatest
role in determining the properties of an
– electron
Chapter 6 (continued)
• What is the periodic trend for:
– Atomic radius
• Increase Top to Bottom and Right to Left
– Ionic radius
• Increase Top to Bottom and Right to Left for both
Cations and Anions
– Ionization energy
• Increase Bottom to Top and Left to Right
– Electronegativity
• Increase Bottom to Top and Left to Right
Chapter 6 (continued)
• What is the energy required to remove an
electron from an atom in the gaseous state
– Ionization Energy
• Which element would have the lowest first
ionization energy?
– Na, Mg, Al, Si, S, or Cl
Which element has the largest atomic radius?
H, Li, Na, K, Rb, Cs or Fr
Li, Be, B, C, N, O, or F
Chapter 6 (continued)
• Which element has the highest
electronegativity value on the periodic table?
– Fluorine
• Define ion, cation, and anion.
– Ion: charged atom
– Cation: positively charged atom
– Anion: negatively charged atom
Chapter 6 (continued)
• What are the ionic charges for the
representative elements?
+1, +2, +3, +/- 4, -3, -2, -1
• Write the charges for each one of these ions
as a superscript
Li +1
Chapter 7
• How many valence electrons would these
atoms have?
– Li:
– N:
– O:
– Cl:
1 valence electron
5 valence electrons
6 valence electrons
7 valence electrons
Chapter 7 (continued)
• What is an ionic compound?
– A compound composed of cations and anions
– Held together by ionic bonds
• What are some characteristics of ionic
– Can conduct electric current when dissolved or in
molten state
– Has a high melting point
– Crystalline solid
– Brittle
Chapter 7 (continued)
• What is the electron configurations of the
following ions:
– N3– Al3+
– Cl1– Ca2+
1s2 2s2 2p6
1s2 2s2 2p6
1s2 2s2 2p6 3s2 3p6
1s2 2s2 2p6 3s2 3p6
Chapter 7 (continued)
• What is a metallic bond?
– The attraction of metal ions to mobile electrons
• What are some properties of metals?
– Conductive (heat and electricity)
– Malleable
– Ductile
– Luster
• Why do metals have these properties
– They have mobile valence electrons
Chapter 8
• In a single covalent bond, how would two
atoms still achieve a stable noble-gas electron
– Two atoms share two electrons
• How many electrons are required to create a
single, double and triple covalent bond?
– 2 shared electrons makes a single covalent bond
– 4 shared electrons makes a double covalent bond
– 6 shared electrons makes a triple covalent bond
Chapter 8 (continued)
• What is an unshared pair of electrons
– Pair of valence electrons that does not participate
in bonding
• According to VSEPR theory, why do molecules
adjust their shapes?
– to keep pairs of valence electrons as far apart as
Chapter 8 (continued)
• What is a polar covalent bond? What is a nonpolar covalent bond? How do you determine
which type of bond would form?
– Polar Covalent bonds are the result of uneven
sharing of electrons
• Based on the range of electronegativity values for the
two atoms
• Those atoms closest to Fluorine will acquire the most
polar charge
– Non-polar covalent bonds are the result of an even
sharing of electrons
Chapter 8 ( continued)
• Which type of solid has the strongest bond
structure and therefore the highest melting
– Ionic Solids
– Metallic Solids
– Network Solids

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