Download Chemistry: Fall Final Review 08

Name__________________________
Date _____________ Period _______
Chemistry: Fall Final Review 08
Safety & Scientific Method
1) What are the safety rules for goggles, chemical safety, lab procedure, fire safety, diluting acids, etc.
Review general safety rules; remember to dilute an acid always pour acid into water
2) Draw a picture of the biological, toxic, open flame, extreme temperature, and sharp object safety
symbols.
Biohazard
Toxic
Open Flame
Extreme temp
Sharp object
3) Give the location of safety devices such as the vent hood, eyewash, fire blanket, fire extinguisher, safety
shower, etc. in the laboratory room.
4) List and describe the steps of the scientific method.
Observations
Hypothesis
Experiment – will have a respondent (dependent) variable and an independent (manipulated) variable
Analyze and conclude results
Name__________________________
Date _____________ Period _______
5) Define qualitative and quantitative.
Qualitative – deals with the type of substances
Quantitative – deals with the amount of substances
Chapter 1 & 3 – Introduction to chemistry
6) Define chemistry.
Chemistry – the study of the structure and composition of matter
7) Name and define the five types of chemistry.
5 types of chemistry:
Organic chemistry – study of substances containing carbon
Inorganic chemistry - study of substances not containing carbon
Analytical chemistry – study of the composition of substances
Physical chemistry – study of the physical behavior of chemicals
Biochemistry – study of the chemistry of living things
8) Define mass and matter.
Mass – the amount of matter in an object
Matter – anything with mass and takes up space
9) Define and describe the states of matter.
Solid – definite shape and volume
Liquid – definite volume
Gas – no definite shape or volume
Plasma
10) What is a physical property? Give examples.
Physical properties can be observed using your senses (color, melting point, state of matter, etc.)
11) What is a chemical change? Give examples.
Chemical changes occur when something new is produced (burning, tarnishing, reacting, etc.)
Name__________________________
Date _____________ Period _______
12) What are the differences between elements and compounds?
Elements are made of only one type of atom; Compounds are made of more than one type of atom
13) What is a mixture?
Physical blend of 2 or more pure substances in any proportion in which each substance retains its
individual properties
14) What are the differences between a homogeneous and heterogeneous mixture?
Homogeneous mixture – you can’t see the different parts
Heterogeneous mixture – you can see the different parts
15) What are the four evidences that a chemical reaction has taken place?
Bubbles, heat, color change, precipitate
16) List & describe the separation techniques of mixtures.
Distillation – separation of two liquids (based on boiling points)
Filtration – separation of a solid and liquid
Crystallization – forms a solid structure
Chromatography – based on charge attraction to another material
17) Give the procedure for separating salt and sand and iron filings.
Use a magnet to separate the iron; dissolve the salt in water and filter the sand
18) Define and give an example of the law of conservation of mass.
4 C3H5(NO3)3  4 CO2 + 6 N2 + 10 H2O + 9 O2
19) Define and give an example of the law of definite proportions?
Regardless of the amount the elements will always be in the same ratio
Ex. Water in a glass and water in a pool – H2O – ratio doesn’t change
20) Define and give an example of the law of multiple proportions?
Two elements can combine in different ratios in order to give different compounds
NO3 – nitrate
Name__________________________
Date _____________ Period _______
NO2 – nitrite
NO – hyponitrite (nitric oxide)
21) What is the difference between mass and weight.
Mass – measurement of matter
Weight – measurement of the force of gravity upon mass
Chapter 2 – Dimensional Analysis & Measurement
22) What would 38 000 000 000 be in scientific notation?
3.8 x 1010
23) Subtract 1.2x107 – 9.1x106.
2.9 x 106
24) (1x103) x (5x10-1) =
5 x 102
25) Rules for Significant figures:
26) Know the rules for significant figures. How many are in: 1.70
Significant figures =
3
0.00550
3
20,000
1
10900
3
27) Define precision and accuracy. Give an example of a measurement that is precise but not accurate.
Precise measurements are close to each other
Accurate measurements are close to the actual value
28) % error – question – removed!!!
29) How many liters are in 12.7 kL?
12,700 Liters
30) Convert 22.5 L to mL.
22,500 mL
Name__________________________
Date _____________ Period _______
31) Use dimensional analysis to convert 22.5 km/hr into m/sec.
22.5 km x 1000 m x 1 hour x 1 min = 6.25 m/s
1 hour
1 km
60 min
60 sec
32) Give an example of everyday liquids (water, oil, Rubbing [isopropyl] alcohol) with different densities.
How would they be arranged if poured together?
Water is more dense than oil, therefore oil floats on top of water
33) What is the density in g/ml of a substance that has a mass of 3.99g and a volume of 0.88ml?
D=m
V
D = 3.99 g
D = 4.53 g/mL
0.88 mL
34) What is the volume of a liquid that has a density of 2.25g/ml and a mass of 18.7g?
V=m
D
V = 18.7 g
V = 8.31 mL
2.25 g/mL
35) What are the temperature conversions from Celsius to Kelvin?
°C + 273 = K
Chapter 4 – Atomic Structure
36) What was Democritus’ contribution to the atomic structure?
He came up with the idea of atoms.
37) Give all of Dalton’s laws.
Elements are composed of tiny indivisible particles called atoms; all atoms of the same element are
identical; atoms can combine in whole number ratios to form compounds
38) Describe Rutherford’s gold foil experiment and what it was used to determine.
Rutherford shot particles through a thin piece of gold foil; some of these particles occasionally
bounced back giving Rutherford the idea that they were bouncing off of a dense area that he called
the nucleus
Name__________________________
Date _____________ Period _______
39) What are the charges and locations of a proton, neutron and electron?
Proton – positive
Neutron – neutral
Electron - negative
40) What information does the atomic number give you?
The atomic number tells you the number of protons or the number of electrons in a neutral atom
41) What is an isotope?
Isotopes of an element have a different mass number and a different number of neutrons
42) Define atomic mass.
The weighted average mass of the isotopes of that element
43) Define mass number.
Number of neutrons + protons
44) How do you find the number of neutrons in an atom? How many neutrons are in carbon-14?
14 – 6 = 8 neutrons in Carbon 14
# neutrons = mass # - atomic #
45) What are the different atomic models? Which is the most current.
Plum Pudding Model – only showed electrons
Rutherford’s Model – included a nucleus
Bohr Model – showed electrons in energy levels
Quantum Mechanical Model – shows the probable locations of electrons using the solutions from
Schrodinger’s equations
46) Define and give an example of Alpha, Beta, and Gamma decay.
1. Alpha (α) – an alpha particle is 2 protons and 2 neutrons (42α)
206
4
Ex. 21084Po
82Pb +
2α
2. Beta (β) – a beta particle is the lost of an electron (when a neutron is converted to a proton) ( 0-1e-)
14
0
Ex. 146C
7N +
-1e3. Gamma (γ) – a gamma particle is a photon at high frequency
31
Ex. 3115P*
15P + γ
Chapter 5 - Electrons
47) What formula for the speed of light?
c = λƒ
Name__________________________
Date _____________ Period _______
48) What is the formula for the Energy of a photon?
E = hƒ
49) Define photon.
A particle of electromagnetic radiation with no mass that carries a quantum of energy
50) Define atomic emission spectra. What does it tell us about the atoms of elements?
A set of frequencies of electromagnetic waves given off by atoms of an element
51) Define Heisenberg uncertainty principle.
It is not possible to know precisely both the velocity and the position of a particle at the same time.
52) Define Aufbau’s principle.
Electrons enter orbitals of the lowest energy first.
53) Define Pauli exclusion principle.
No 2 electrons in the same atom can have the same set of 4 quantum numbers
54) Define Hund’s Rule.
Electrons in the same subshell occupy available orbitals singly before pairing up
(1 e- per orbital till each orbital has one then begin putting a 2nd e- in each orbital)
55) Define Valence electrons.
The electrons in the outermost orbital
Help determine an atoms properties
56) Define Octet rule.
Atoms lose or gain or share electrons in order to aquire a full set of 8 valence electrons.
(first shell has 2 valence electrons but all others have 8)
57) How many electrons can each orbital hold? How many electrons can the s, p, d and f sublevels hold?
Each orbital can only hold 2 electrons
S sublevel – has one orbital = 2 electrons total
P sublevel – has 3 orbitals = 6 electrons total
D sublevel – has 5 orbitals = 10 electrons total
58) How can you easily locate the energy level by using periodic table?
Locate the row of the element
Name__________________________
Date _____________ Period _______
59) Draw the orbital notation for chlorine.
60) What is the electron configuration of lead.
1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p2
61) In 2s1, what do the 2, the s and the 1 represent?
The 2 represents the energy level and the 1 represents one electron.
62) How many valence electrons are in N, Cl and Na?
N has 5 valence e- , Cl has 7 valence e- , Na has 1 valence e-
(same as the group number)
63) What are the 4 quantum numbers of the last electron in Cr, S, I, Mg? OMIT!!!
Cr – (3d4) - 3,2,1, ½
S – (3p4) – 3,1,-1, - ½
I – (5p5) – 5,1,0, - ½
Mg – (3s2) – 3,0,0, - ½
64) Write out a stable electron configuration of Copper.
[Ar] 4s23d9 - normal
[Ar] 4s13d10 - [Ar] 4s03d10 – CU+1
[Ar] 4s23d9 - [Ar] 4s03d9 – CU+2
Chapter 6 – Periodic Table
65) How was Mendeleev’s periodic table arranged?
By increasing atomic mass
66) How was Mosley’s periodic table arranged?
By increasing atomic number
67) What is the group 1A elements called? Group 2A? Group 7A? Group 8A? What kind of elements are in
groups 3-12?
Group 1A – Alkali Metals
Group 2A – Alkaline Earth Metals
Group 17 (Group 7 A) – Halogens
Group 18 (Group 8 A) – Noble Gases
Groups 3-12 – Transition Elements
Name__________________________
Date _____________ Period _______
68) How are elements grouped on the periodic table?
Valence electrons
69) What is the trend for atomic radius? What element has the largest atomic radius?
Atomic radius increases from top to bottom on the periodic table and it increases from right to left.
The largest atomic radius is Francium
70) Place in increasing radius - carbon, a carbon cation or a carbon anion?
Carbon cation < Carbon atom < carbon anion
71) What is ionization energy? What is it’s trend?
Ionization energy is the amount of energy it takes to remove an electron from an atom. Ionization
energy increases from bottom to top and it also increases from left to right on the periodic table.
Helium has the highest ionization energy.
72) What is the trend in electronegativity? Where are the elements with the smallest electronegativity?
Electronegativity increases from left to right (not including the noble gases) and it increases from top
to bottom. The smallest electronegativity elements are the noble gases; the biggest is Fluorine.
Chapter 8 – Ionic Bonding
73) Define Ionic Bond.
Complete transfer of electrons from one atom to another
Bond between ions (cation and anion)
74) Define cation and anion.
Cation – a positive ion resulting from the loss of electrons
Anion – a negative ion resulting from the gain of electrons
75) Gaining three electrons gives what overall charge?
Negative 3 charge
Name__________________________
Date _____________ Period _______
76) Draw the Lewis dot structure for N, K, & Cl-.
N
K
Cl-
77) Name the following compounds:
CaCl2 - Calcium Chloride
CuNO3 - Copper (I) Nitrate
CCl4 - Carbon Tetrachloride
Ba(C2H3O2)2 - Barium Acetate
Si3H8. - Trisilicon Octahydride
78) Write the formula for the following:
zinc bromide ZnBr
dinitrogen monoxide - N2O
lead (IV) oxide - Pb2O
copper (I) carbonate - Cu2CO3
calcium fluoride. – CaF2
79) When is an ionic compound able to conduct electricity and why?
An ionic compound can conduct electricity when it is in solution (in water). This allows the ions to
float around freely and conduct current.
80) What properties of metals are explained by metallic bonding theory?
Metals are ductile and malleable due to their metallic bonding and the sea of free floating electrons.
81) Define and give an example of an alloy.
Mixture of 2 metals.
Steel
Name__________________________
Date _____________ Period _______
Chapter 9 – Covalent Bonding
82) Give examples of pairs of elements that would form ionic bonds, nonpolar bonds and polar bonds.
Nonpolar bonds are between elements with less than a 0.4 difference in electronegativity (ex: Cl2)
Polar bonds have elements with between a 0.4 and 1.7 difference in electronegativity (ex: CCl4)
Ionic bonds have more than a 1.7 difference in electronegativity (NaCl)
83) What are sigma and pi bonds?
Sigma bond – formed by end to end overlap of an s orbital
Pi bond – formed by side to side overlap of p orbitals
84) Compare and contrast ionic compounds to covalent compounds?
Ionic compounds transfer electrons between atoms. Covalent compounds share electrons.
85) What is VSEPR?
Valence Shell Electron Pair Repulsion
Electron move as far away from each other as possible
86) How do you determine Molecular shape of a molecule?
By the shared and unshared valence electrons (atoms attached and the lone pairs)
87) Draw the Lewis structure and list the molecular shape for these molecules:
SeF2
- Bent
CO2 - Linear
CH4 - Tetrahedral
NH3 – Trigonal pyramidal
HCl - Linear
CH3CH2OH – central atom - tetrahedral