Study Guide for

NYS Regents Chemistry

Midterm Examination

Table of Contents:

Topic Page

1. Scientific Method ……………………………………………………………….. 1

2. Significant Figures ……………………………………………………………… 2

3. Scientific Notation ……………………………………………………………… 4

4. Accuracy and Precision ………………………………………………………… 4

5. Graphs …………………………………………………………………………... 4

6. Density ………………………………………………………………………….. 4

7. Error Calculations ………………………………………………………………. 5

8. Unit Changes ……………………………………………………………………. 5

9. Standard Units …………………………………………………………………… 5

10. Atomic Structure ………………………………………………………………... 6

11. Periodic Table …………………………………………………………………... 9

12. Chemical Formulas ……………………………………………………………... 11

13. Mole Information ……………………………………………………………….. 14

14. Chemical Equations …………………………………………………………….. 16

15. Matter …………………………………………………………………………… 18

16. Energy and Heat ………………………………………………………………… 19

17. Gas Laws ………………………………………………………………………… 21

MHS Regents Chemistry Midterm Review 3rd Edition

1)  Scientific Method – a logical, systematic approach to the solution of a scientific problem.

a)  Make observations. An observation can lead to a question.

b)  Hypothesis – An educated guess based on observed facts. A hypothesis can be revised based upon experimental data.

c)  Controlled Experiments – All factors or variables are held constant while only one variable is changed at a time in order to see the effect of that variable on the experiment.

d)  Data – The results of an experiment, which often include a collection of measurements

e)  Theory – Provides a general explanation for the observations made of many scientists working in different areas of science over a long period of time. Answers the “why” and “how” questions.

f)  Modify Theories if necessary.

g)  Scientific Laws. A Scientific Law simply states a relationship between observed facts. It describes natural events but does not explain why or how they occur.

2)  Significant figures – The digits in a number that represent a quantity actually measured. Significant figures (sig.figs.) indicate the “accuracy” in a number.

a)  Obtaining laboratory measurements with correct number of significant figures (accuracy): use the maximum number of units available – read carefully

b) Rules for deciding the number of significant figures in a measured quantity:

(1) All nonzero digits are significant:
1.234 g has 4 significant figures,
1.2 g has 2 significant figures.
(2) Zeroes between nonzero digits are significant:
1002 kg has 4 significant figures,
3.07 mL has 3 significant figures.
(3) Leading zeros to the left of the first nonzero digits are not significant; such zeroes merely indicate the position of the decimal point:
0.001 oC has only 1 significant figure,
0.012 g has 2 significant figures.
(4) Trailing zeroes that are also to the right of a decimal point in a number are significant:
0.0230 mL has 3 significant figures,
0.20 g has 2 significant figures.
(5) When a number ends in zeroes that are not to the right of a decimal point they are not significant:
190 miles has 2 significant figures,
50,600 calories has significant figures.

c)  Multiplication/Division Rule using Sig.Figs.

i)  The number of SF in the answer cannot be more than the least number of Sig. Figs. used in the calculation e.g. 3.25 x 0.025 = 0.081 (having 2 Sig. Figs. in the answer)

d)  Addition/Subtraction Rule using Sig.Figs.

i)  Line up the decimal places of each number, calculate the answer and round the answer to the smallest decimal place in any of the numbers used.

Question / Answer / Question / Answer
2.54 x 15 x 352 = / 203.0 / 0.20 =
(50500)(2.5)/(15.0) = / (456.03)/((2.67)(301)) =
2.135 + 53.24 + 873.2 = / 485.23 – 32.1 =
85.967 – 0.12 = / 0.0508 – 0.004 =

3)  Scientific Notation (SN)

  1. Change the decimal to SN
  2. 0.00013 to 1.3 x 10-4 : move the decimal point between the first and second significant figure and the power of ten is the number of places moved (right is negative and to left is positive); number smaller than 1.0000 has a negative power of ten.
  3. Change SN into a decimal
  4. 5.4 x 103 to 5400 positive power of ten – move decimal point to the right the number of spaces equal to the power of ten; negative power of ten – move decimal point to the left the number of spaces equal to the power of ten
  5. Preserve the number of significant figures when changing between decimal and SN!

Decimal / # Sig. Figs. / Sci. Notation / Sci. Notation / # Sig. Figs. / Decimal
0.0020300 / 5 / 2.0300 x 10-3 / 2.05 x 105 / 3 / 205,000
  1. Enter SN, 1.3 x 10-4, into your NON-TI83 calculator
  2. Casio calculators usually enter the number given above as “1 . 3 EXP – 4”
  3. TI calculators usually enter the number given above as “1 . 3 (2nd for TI83) EE – 4”

4)  Accuracy compared to Precision

e.  Precision – The reproducibility of a series of measurements

f.  Accuracy – How close a measurement is to the true or accepted value.

5)  Graphs

  1. Horizontal axis (abscissa) – the independent variable
  2. The variable you control
  3. Vertical axis (ordinate) – the dependent variable
  4. The variable you measure as a result of your experiment

6)  Density: Formula can be found on Table t

  1. D = mass / volume
  2. Volume can be calculated by direct measurement: (length) x (width) x (height)
  3. Volume can be calculated indirectly by volume displacement: (volume of liquid with submerged substance) – (volume of liquid without submerged substance)
  4. D = (gram formula mass)/(22.4L) for gases at STP

7)  Error calculations:

  1. Relative (percent) error: (Be careful to divide by the accepted value and not the measured one.)

measured value – accepted value

Percent error = ------x 100

accepted value

8)  Unit Changes Table C of Reference Tables

  1. Using dimensional analysis (Factor Label Method)
  2. Starting unit in numerator (over 1), then multiply by ratio of units such that starting unit is in denominator and new unit is in the numerator. Continue until desired unit is reached.
  3. K H D (m) D C M
  4. Kids Hate Doing Math During Certain Months
  5. Kilo Hecto Deca (unit) Deci Centi Milli
  6. k h da d c m
  7. change units by moving from starting unit to final unit and move decimal point in the same direction by the same number of moves.

9)  Standard Units Table D of Reference Tables

  1. Standard International Units (SI units)
  2. Time: second
  3. Temperature: K (Kelvin)
  4. Mass: kg (kilogram)
  5. Length: m (meter)
  6. Amount of substance: mole

10) Atomic Structure

  1. Models of the atom developed over a long time.
  2. Dalton: solid atom as a ball and is the smallest indivisible unit of matter
  3. Thomson: discovered the electron; used cathode ray tube, small negative particles (electrons) came out of all matter, Plum Pudding Model
  4. Rutherford: discovered the nucleus; Gold foil experiment shot alpha particles at a gold foil and almost all particles went straight through only a few deflected away; Conclusions: atoms are made up of mostly empty space with a small, dense, positively charged center. The electrons were somewhere outside the nucleus in the “mostly empty space.”
  1. Bohr: electrons must be in energy levels at specific distances from the nucleus and never between these levels; concluded this from examining the bright line spectrum of hydrogen atoms. Each energy level has a specific energy. The further the level is away from the nucleus the greater the energy of the electrons in it.
  2. Bright line spectrum: When an electron in an atom gains just the right amount of energy, from an outside source, electron can shift to a higher energy state (excited state). However, the excited state atom is energetically unstable and the electrons will return to the ground state – giving off the absorbed energy in the form of light. The amount of energy in the released light (wavelength) depends on how many energy levels the electron jumps back, how many other electrons are around and also the charge of the nucleus. Every element has a different number of electrons and protons and therefore will produce different light energies. The pattern of light colors produced by an element is its bright line spectrum and is unique to that element.
  3. Although atomic line spectra were known before Bohr proposed his model of the hydrogen atom, Bohr was able to apply mathematics to his model and was able to account for each line in the visible spectrum of hydrogen. The Bohr model failed to explain the energies absorbed or emitted by atoms with more than one electron (hydrogen)
  1. Electron Cloud (Quantum Model)(Modern Model)(Wave-mechanical Model): electrons arranged around atomic nucleus in bands or regions likely to contain electrons
  2. Orbitals are regions around the nucleus that electrons are most likely to be found.
  3. Protons: particle in nucleus; +1 charge; mass = 1 amu (atomic mass unit)
  4. Neutron: particle in nucleus; no charge; mass = 1 amu
  5. Electron: particle located outside the nucleus; -1 charge; negligible mass (amu)
  6. 1 amu (atomic mass unit) = mass of one 12C atom.
  1. Atomic Number: defines an element; number of protons in the nucleus of an atom e.g. 7N nitrogen has 7 protons in its nucleus; also is equal to the number of electrons in any uncharged atom.
  2. Mass Number: (number of protons) + (number of neutrons) in the nucleus of an atom. Mass Number = number of nucleons
  3. Number of neutrons in a nucleus calculated by (Mass Number)-(Atomic Number)
  1. Isotope: atoms of a same element (atomic number) but have a different mass number; different number of neutrons e.g. 12C and 14C; also indicated as C-12 and C-14
  1. Average atomic mass: the weighted average mass of all naturally occurring isotopes, shown on periodic table as mass e.g. 14.0067N;Average atomic mass = (decimal of percent abundance)(mass of that isotope)+(decimal of percent abundance)(mass of that isotope)+(decimal of percent abundance)(mass of that isotope) etc. for each isotope. (NOTE: if isotope masses are not given then use the isotope mass numbers instead) – Use the link below for a tutorial and video:

http://www.kentchemistry.com/links/AtomicStructure/atomicmasscalc.htm

Example: The element copper has naturally occurring isotopes with mass numbers of 63 and 65. The relative abundance and atomic masses are 69.2% for mass = 62.93amu, and 30.8% for mass = 64.93 amu. Calculate the average atomic mass of copper.

(69.2 x 62.93) + (30.8 x 64.93) =

100

4354.76 + 1999.84 = 63.55

100

  1. Nucleons: particles in the nucleus are protons and neutrons

z.  Ions

i. Negative ions (anions) are formed when atoms gain electrons; larger than atom

Atomic radius < Ionic radius

ii.  Positive ions (cations) are formed when atoms lose electrons; smaller than atom

Atomic Radius > Ionic radius

11)  Periodic Table

  1. Elements were at first arranged according to mass
  2. Today elements are listed in order of atomic number (number of protons in the nucleus)
  3. Elements in the same row (period) have the same number of occupied principal energy levels (period number also equals number of outermost principal energy level containing electrons), but have different chemical and physical properties
  4. Elements in the same group (column) have the same number of electrons in their outermost level and therefore have similar chemical properties
  5. Group 1 “Alkali Metals”; Group 2 “Alkaline Earth Metals”; Groups 3-11 “Transition Metals”; Group 17 “The Halogens”; Group 18 “Noble Gases”
  6. Valence shell: the outermost level containing electrons

i.  Lewis Dot Structure: Is the atomic symbol with its valence electrons drawn around it.

1.  Atoms (example): and ּּ

2.  Ions (examples): [Ca]+2 ; [::] -2

  1. Kernel – part of atom NOT including valence electrons
  2. Three types of elements: metals, non-metals and metalloids
  3. Metals and non-metals separated by “staircase” beginning at Group 13
  4. Metals to the left of the “staircase” (except H) (most elements are metals)
  5. Non-metals to the right of the “staircase” (including H)
  6. Properties of Metals:
  7. Are mostly solids (one liquid, Hg)
  8. Lose electrons easily (low ionization energy)
  9. Don’t want to gain more electrons (low electronegativity)
  10. Good conductors of heat and electricity (due to mobile electrons)
  11. Ductile – ability to be drawn into wires
  12. Malleable – ability to be rolled or hammered into thin sheets
  13. Luster – glossy finish, reflect light
  14. Want to give away electrons to form positive ions (cations) with smaller radii
  15. Metallic properties are more pronounced with increasing atomic size and fewer valence electrons.
  1. Transition elements are elements in groups 3 through 11
  2. These elements generally have a color when they are ions (either in a compound or dissolved in water).
  3. Properties of Non-metals:
  4. Are mostly gases (some solids, one liquid, Br2)
  5. Want to gain electrons and form negative ions (anions) with larger radii (high electronegativity)
  6. Don’t want to give up electrons (high ionization energy)
  7. Poor conductors of heat and electricity
  8. Solids tend to be brittle
  9. Solids generally have a dull finish
  10. Metalloids: 6 elements on the “staircase” except Al, At, and Po which are metals; have properties of both metals and non-metals (semimetals or semiconductors)
  1. All elements with atomic numbers greater than ‘83’ have no known stable isotopes. They are all radioactive.
  2. Diatomic elements: Br I N Cl H O F – Bromine, Iodine, Nitrogen, Chlorine, Hydrogen, Oxygen, and Fluorine.
  1. Gas, Liquid, and Solid Elements
  2. 11 Gases at STP: Fluorine, oxygen, hydrogen, chlorine, nitrogen (last 5 of Br I N Cl H O F) plus all 6 of the Noble Gases
  3. 2 Liquid elements at STP: Bromine and Mercury (Hg)
  4. ALL OTHER ELEMENTS ARE SOLIDS at STP
  5. Atomic radius
  6. Increase size going down any group (new electron shells and increased nuclear shielding)
  7. Decrease size going left to right across any Period (row) (increased nuclear charge)
  8. Electronegativity: the “pull” atoms of an element have to gain another electron
  9. First Ionization Energy: the energy necessary to remove the first valence electron from an atom in the gaseous state.
  10. Fluorine is the most electronegative element; most active non-metal
  11. Francium is the least electronegative element; most active metal
  1. Highly reactive elements – found only as compounds in nature
  2. Group 1: Alkali Metals
  3. Group 2: Alkaline Earth Metals
  4. Group 17: Halogens (has elements in all three phases)
  5. The most reactive of these can only be purified by electrolysis
  6. Group 18: Noble Gases (Monatomic Molecules)
  7. Generally not reactive – already have the stable octet valence electron configuration
  1. Allotropes: one or more molecular forms of the same element in the same state. Examples: carbon as diamond, graphite and coal; oxygen as O2 or ozone, O3
  2. Ground state electron configuration: when electrons of an atom are occupying the lowest possible energy levels, which is the electron configurations given on the Periodic Table. For example: Mg is 2-8-2
  3. Excited state electron configurations: when an electron(s) are in a higher energy level than the lowest possible energy levels. Example: excited Mg 2-7-3

12) Chemical Formulas