1. Chemistry Review

2. Matter and Change
Chm.1.1 – Analyze the structure of atoms, isotopes, and ions.
Chm.1.2 – Understand the bonding that occurs in simple compounds in terms of bond type, strength, and properties.
Chm.1.3 – Understand the physical and chemical properties of atoms based on their position on the periodic table.

3. Chm.1.1.1 – Analyze the structure of atoms, isotopes, and ions.
Characterize protons, neutrons, electrons by location, relative charge, relative mass (p=1, n=1, e=1/2000).
Use symbols: A = mass number, Z = atomic number
Use notation for writing isotope symbols: 23592U or U-235.
Identify isotope using mass number and atomic number and relate to number of protons, neutrons and electrons.
Differentiate average atomic mass of an element from the actual isotopic mass and mass number of specific isotopes.

4. Energetic particles that move in all directions around the nucleus of an atom are called:
Neutrons.
Protons.
Elements.
Electrons.

5. An ion with a net charge of +2 loses 5 electrons. What is its charge?7+ 3+ 3- 7-

6. What is the net charge of an ion with 4 protons and 6 electrons?
24+
10+ 2-
Neutral

7. What is the total number of subatomic particles in
an atom of potassium-39? 19 20 39 78

8. The number of protons in the nucleus of an atom of a specific element is the same as that element’s:
Atomic mass
Energy levels
Atomic number
Neutrons

9. The number 80 in the name of bromine-80 represents
The atomic number
The mass number
The sum of protons and electrons
None of these

10. The table gives the natural percent abundance of the stable isotopes of sulfur. Based on the data, what is the reasonable estimate for the average atomic mass of sulfur?
33.7
35.0
32.1
34.0
None
Isotope
Natural % Abundance
Mass (amu)
Sulfur-32
95.022
31.972
Sulfur-33
0.76
32.971
Sulfur-34
4.22
33.967
Sulfur-36
0.014
35.967

11. Chm.1.1.2 – Analyze an atom in terms of the location of electrons.
Analyze diagrams related to the Bohr model of the hydrogen atom in terms of allowed, discrete energy levels in the emission spectrum.
Describe the electron cloud of the atom in terms of a probability model.
Relate electron configurations of atoms to the Bohr and electron cloud models.

12. The quantum mechanical model of the atom:
Is concerned with the probability of finding an electron in a certain position.
Was proposed by Neils Bohr.
Defines the exact path of an electron around the nucleus.
Has many analogies in the visible world.

13. What best describes how the Bohr model differs from the quantum mechanical model of the atom?
The quantum model does not define an exact pattern the electrons take around the nucleus.
The quantum model is based on fixed energy levels of electrons.
The Bohr model restricts the energy of electrons to certain values.
The Bohr model uses probability to determine the location of finding an electron around the nucleus.

14. Which best describes Bohr’s model of a hydrogen atom?
The electron bound in a circular orbit around the nucleus
A large, dense nucleus surrounded by atoms in different orbits
A nucleus surrounded by electrons in specific energy levels
All the particles mixed together like “plum pudding”

15. Chm – Explain the emission of electromagnetic radiation in spectral form in terms of the Bohr model.
Understand that energy exists in discrete units called quanta.
Describe the concepts of excited and ground state of electrons in the atom:
Gaining energy results in the electron moving from its ground state to a higher energy level.
When the electron moves to a lower energy level, it releases the energy difference in the two levels as electromagnetic radiation (emissions spectrum).

16. Chm – Explain the emission of electromagnetic radiation in spectral form in terms of the Bohr model
Understand that electromagnetic radiation is given off as photons.
Use the “Bohr Model for Hydrogen Atom” and “Electromagnetic Spectrum” diagrams from the Reference Tables to relate color, frequency, and wavelength of the light emitted to the energy of the photon.
Understand the inverse relationship between wavelength and frequency, and the direct relationship between energy and frequency.

17. Chm – Explain the emission of electromagnetic radiation in spectral form in terms of the Bohr model
Explain that Niles Bohr produced a model of the hydrogen atom based on experimental observations. This model indicated that:
An electron circles the nucleus only in fixed energy ranges called orbit
An electron can neither gain or lose energy inside this orbit, but could move up or down to another orbit
That the lowest energy orbit is closest to the nucleus

18. The frequency and wavelength of all waves are
Directly related.
Inversely related.
Unrelated
Equal

19. Once the electron in a hydrogen atom absorbs a quantum of energy, it
Is now in its ground state
Is now in its excited state
Has released a photon
Follows an exact path around the nucleus

20. Visible light is part of the electromagnetic spectrum
Visible light is part of the electromagnetic spectrum. If a light wave had a wavelength of x 10-7 nm, what color would it appear?
Violet
Blue
Green
Yellow

21. According to the Bohr Model for a hydrogen atom, which change in energy level would emit visible light?
n = 2 to n = 1
n = 2 to n = 3
n = 4 to n = 2
n = 4 to n = 6

22. According to the Bohr Model for a hydrogen atom, what wavelength of light would be emitted if an electrons moved from n=3 energy level to its ground state?
103 nm
434 nm
656 nm
1094 nm

23. FLAME TEST
Metal
Flame Color
Stronium
Scarlett
Calcium
Orange
Copper
Blue-Green
Potassium
Purple
Sodium
Yellow
A student performs a flame test on an unknown solution. The unknown solution causes the flame to turn orange. Which element is most likely contained in the solution?
Sodium
Copper
Potassium
Calcium

24. According to the Bohr model for a hydrogen atom, what wavelength of light would be emitted when an electron jumps from n=2 to its ground state?
486 nm
122 nm
102 nm
97 nm

25. Chm.1.1.4 – Explain the process of radioactive decay using nuclear equations and half-life.
Use the symbols for and distinguish between alpha and beta nuclear particles, and gamma radiation , including relative mass.
Use shorthand notation for particles involved in nuclear equation to balance and solve for unknowns.
Compare the penetrating ability of alpha, beta, and gamma radiation.

26. Chm.1.1.4 – Explain the process of radioactive decay using nuclear equations and half-life.
Describe nuclear decay, including:
Decay as a random event, independent of other energy influences
Using symbols to represent simple balanced decay equations
Simple half-life calculations
Compare radioactive decay with fission and fusion.

27. Which type of ionizing radiation can be blocked by clothing?
Alpha particle
Gamma radiation
X-rays
Beta particle

28. If an isotope undergoes beta emission
The mass number changes.
The atomic number changes.
The atomic number remains the same.
The number of neutrons remains the same.

29. Which of the following types of radiation has no mass and no charge?
Alpha
Beta
Gamma
positron

30. When Rn-222 undergoes decay to become Po- 218, it emits
An alpha particle
A beta particle
Gamma radiation
X-rays

31. Which nuclear equation shows the radioactive process for beta emission by argon-37?

32. After 252 days, a 48-g sample of scandium-42 contains only 6
After 252 days, a 48-g sample of scandium-42 contains only 6.0-g of the isotope. What is the half-life of scandium?
84 days
42 days
32 days
28 days

33. When nuclear fission occurs
Two nuclei combine to produce a heavier nucleus
The chain reaction that results cannot be controlled
It is a spontaneous reaction
It must be initiated by bombardment with neutrons

34. Which of the following has a +2 charge?
Alpha particle
Beta particle
Neutrino
Gamma ray

35. Chm.1.2.1 – Compare (qualitatively) the relative strengths of ionic, covalent, and metallic bonds.
Describe metallic bonds :”metal ions plus ‘sea’ of mobile electrons”.
Describe how ions are formed and which arrangements are stable (filled sub-levels).
Appropriately use the term cation as a positively charged ion and anion as negatively charged ion.
Predict ionic charges for representative elements based on valence electrons.
Apply the concept that sharing electrons form a covalent compound that is a stable arrangement.
Draw Lewis structure for simple compounds and diatomic elements indicating single, double or triple bonds.

36. Why does a cation have a positive charge?
It has lost valence electrons.
It has gained valence electrons.
It has an ionic bond.
It has a metallic bond.

37. How many electrons must an atom of strontium lose to have to gain a noble gas electrons configuration?
None 1 2 3

38. Which group of elements from the periodic table gain one electron when they become ions?

39. How many valence electrons do the elements in Group 6A have?8 6 2

40. In general, metals react by
Losing valence electrons
Gaining valence electrons
Sharing valence electrons
Sometimes gaining and sometimes losing valence electrons

41. Which of the following is a diatomic molecule containing a triple covalent bond?
N2O
Br2
N2O2

42. Chm.1.2.2 – Infer the type of bond and chemical formula formed between atoms.
Determine that a bond is predominantly ionic by the location of the atoms on the Periodic Table (metals combined with nonmetals) or when ∆EN>1.7.
Determine that a bond is predominantly covalent by the location of the atoms on the Periodic Table (nonmetals combined with nonmetals) or when ∆EN<1.7.
Predict chemical formulas of compounds using Lewis structures.

43. A covelent bond forms
When an element becomes a gas
When atoms share electrons
Between metals and nonmetals
When electrons are transferred from one atom to another

44. Which of the following compounds contain ionic bonds?NO
FeO
CH4

45. Which of the following is a true statement?
When the electronegativelity difference between two atoms is greater than 1.7, the atoms form a covalent bond.
The size of th electronegativity difference between two atoms has no bearing on the type of bond that the two atoms may form.
As the electronegativity difference between two atoms increases, the polarity of the bond increases.
If the electronegativity difference between two atoms is 0.3, the atoms have a weak polar bond.

46. Which pair of atoms is most likely to form an ionic
bond?
Mg and S
C and O
Br and S
Ca and O

47. Chm.1.2.3 – Compare inter- and intra- particle forces.
Explain why intermolecular forces are weaker than ionic, covalent or metallic bonds.
Explain why hydrogen bonds are stronger than dipole-dipole forces which are stronger than dispersion forces.

48. Chm.1.2.3 – Compare inter- and intra- particle forces.
Describe intermolecular forces for molecular compounds.
H-bond as attraction between molecules when H is bonded to O, N, or F.
Dipole-dipole attractions between polar molecules.
London dispersion forces (electrons of one molecules attracted to nucleus of anotehr molecule – i.e. liquefied inert gases
Relative strengths – (H>dipole>London/van der Waals)

49. Which molecules has the strongest forces of attraction between molecules?
HBr HF
HCl

50. Chm.1.2.4 – Interpret the name and formula of compounds using IUPAC convetion.
Write binary compounds of two nonmetals (use Greek prefixes)
Write binary compounds of metal/nonmetal.
Write ternary compounds (polyatomic ions) using the polyatomic ions on the reference table.
Write, with charges, these polyatomic ions: nitrate, sulfate, carbonate, acetate, and ammonium.
Know the names and formulas for these common laboratory acids: HCl, HNO3, H2SO4, HC2H3O2 (or CH3COOH)

51. Chm.1.2.5 – Compare the properties of ionic, covalent, metallic, and network compounds.
Apply VSEPR for electron pair geometries.
Describe bond polarity. Polar/nonpolar molecules and solubility.

52. What is the correct formula for sulfur trioxide?
S3O3
SO3 SO
S3O

53. If you see an –ide at the end of a chemical name, what can you assume?
It is a binary compound.
It is an acid.
It has a neutral charge.
It is a polyatomic ion.

54. Which of the following represents the compound formula from these two ions: silver and hydroxide?
Ag2(OH)2
AgOH
Ag(OH)2
2Ag(OH)2

55. What is the formula for aluminum oxide?
AlO3
Al2O3
Al3O2
Al2O

56. What is the name for the compound CuCl2?
Copper (I) chloride
Copper (II) chloride
Copper (I) chlorine
Copper (II) chlorine

57. What is the formula for the compound FeCO3?
Iron carbon oxide
Iron (I) carbide
Iron (I) carbonate
Iron (II) carbonate

58. What is the name for the compound N2O3?
Nitrogen oxide
Nitrous oxide
Nitrogen (II) oxide
Dinitrogen trioxide

59. Chm.1.2.5 – Compare the properties of ionic, covalent, metallic, and network compounds.
Explain how ionic bonding in compounds determines their characteristics: high MP, high BP, brittle, high electrical conductivity either in molten state or in aqueous solution.
Explain how covalent bonding in compounds determines their characterics: low MP, low BP, poor electrical conductivity, polar nature.
Explain how metallic bonding determines the characteristics of metals: high MP, high BP, high conductivity, malleability, ductillity, luster

60. Using VSEPR theory, predict the geometry around the central atom of methane, CH4.
Planar triangular
Tetrahedral
Linear
Bent

61. Test Results
Sample A
Sample B
Melting Point
148
946
Electrical conductivity when dissolved
None
Good
A student performs the same test on two white crysalline solids, A and B. The two results are shown below. Based on the results which statement is true?
Solid A contains only covalent bonds and solid B contains only ionic bonds.
Solid A contains only ionic bonds and solid B contains only covalent bonds.
Both solids contain only ionic bonds.
Both solids contain only covalent bonds.

62. What is the shape of a molecule which has three shared pairs of electrons and no unshared pairs?
Trigonal planar
Tetrahedral
Linear
Bent

63. Which set of characteristics correctly describes an ionic compound?
High melting point, high boiling point, brittle crystals
Low melting point, low boiling point, soft crystals
High melting point, low boiling point, soft crystals
Low melting point, high boiling point, brittle crystals

64. Chm – Classify the components of a periodic table (period, group, metal, metalloid, nonmetal, transition)
Identify groups as vertical columns on the periodic table.
Know that main group elements in the same group have similar properties, the same number of valence electrons, and the same oxidation number.
Summarize that reactivity increases as you go down within a group for metals and decreases for nonmetals.

65. Chm – Classify the components of a periodic table (period, group, metal, metalloid, nonmetal, transition)
Identify periods as horizontal rows on the periodic table.
Identify regions of the periodic table where metals, nonmetals, and metalliods are located.
Classify elements as metals/nonmetals.metallids based on location.
Identify representative (main group) elements as A groups or as groups 1, 2, 13 – 18.
Identify alkali metals, alkaline metals, halogens, and noble gases based on location on the periodic table.
Identify transition elements as B groups or as groups 3 – 12.

66. Which group of metals is so reactive that is members are never found uncombined in nature?
Alkali metals
Halogens
Alkaline earth metals
Noble gases

67. In the periodic table, elements with similar properties are found in the same
Group
Period
Row
Series

68. Which of the following statements about a column of the periodic table is true?
The elements have similar properties.
The elements have a wide range of properties.
The elements have the same atomic number.
The elements have the same atomic mass.