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General Flashcards

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1
Q

The English System of Units consists of:

A
  1. Length: inch (in), foot (ft), yard (yd), mile (mi)
  2. Volume: fluid ounce (oz), cup (c), pint (pt), quart (qt), gallon (gal)
  3. Weight: ounce (oz), pound (lb), ton
  4. Time: second (s), minute (min), hour (h), day (d), year (y)
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2
Q

English System of Units conversion factors:

A
  1. Length:
    12 in = 1 ft; 3 ft = 1 yd; 5280 ft = 1 mi; 1760 yd = 1 mi
  2. Volume:
    2 c = 1 pt; 2 pt = 1 qt; 32 oz = 1 qt; 4 qt = 1 gal
  3. Weight:
    16 oz = 1 lb; 2000 lb = 1 ton
  4. Time:
    60 s = 1 min; 60 min = 1 h; 24 h = 1 d; 365 1/4 d = 1 y
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3
Q

Units of the Metric System:

A
  1. Length: meter (m)
  2. Volume: liter (L)
  3. Mass: gram (g)
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4
Q

Metric System prefixes:

A
  1. femto-
  2. pico-
  3. nano-
  4. micro-
  5. milli-
  6. centi-
  7. deci-
  8. kilo-
  9. mega-
  10. giga-
  11. tera-
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5
Q

Femto-:

A

Symbol: f
Meaning: 10^-15

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6
Q

Pico-:

A

Symbol: p
Meaning: 10^-12

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7
Q

Nano-:

A

Symbol: n
Meaning: 10^-9

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8
Q

Micro-:

A

Symbol: mu
Meaning: 10^-6

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9
Q

Milli-:

A

Symbol: m
Meaning: 10^-3

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10
Q

Centi-:

A

Symbol: c
Meaning: 10^-2

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11
Q

Deci-:

A

Symbol: d
Meaning: 10^-1

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12
Q

Kilo-:

A

Symbol: k
Meaning: 10^3

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13
Q

Mega-:

A

Symbol: M
Meaning: 10^6

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14
Q

Giga-:

A

Symbol: G
Meaning: 10^9

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15
Q

Tera-:

A

Symbol: T
Meaning: 10^12

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16
Q

1 mL =

A

1 cm^3

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17
Q

1 g =

A

1 mL H20 at 4 C

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18
Q

Mass versus Weight:

A

Mass is a measure of the amount of matter in an object, so the mass of an object is constant.
Weight is a measure of the force of attraction of the earth acting on an object. The weight of an object is not constant.

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19
Q

The seven base units for the SI system are:

A
  1. Length: meter (m)
  2. Mass: kilogram (kg)
  3. Time: second (s)
  4. Temperature: kelvin (K)
  5. Electric current: ampere (D)
  6. Amount of substance: mole (mol)
  7. Luminous intensity: candela (cd)
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20
Q

The common derived SI units used in chemistry are:

A
  1. density (kg/m^3)
  2. electric charge: coulomb (C (A s))
  3. electric potential: volt (V (J/C))
  4. energy: joule (J (kg-m^2/s^2))
  5. force: newton (N (kg-m/s^2))
  6. frequency: hertz (Hz (s^-1))
  7. pressure: pascal (Pa (N/m^2))
  8. velocity (speed): meters per second (m/s)
  9. volume: cubic meter (m^3)
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21
Q

The Non-SI units in common use:

A
  1. Volume: liter (L (10^-3 m^3)
  2. Length: angstrom (D (0.1 nm))
  3. Pressure: atmosphere (atm (101.325 kPa))
    torr (mmHg (133.32 Pa))
  4. Energy: electron volt (eV (1.601 x 10^-19 J))
  5. Temperature: degree Celsius (EC (K - 273.15))
  6. Concentration: molarity (M (mol/L))
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22
Q

Unit conversions for length:

A

1 m = 1.094 yd

1 yd = 0.9144 m

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23
Q

Unit conversions for volume:

A

1 L = 1.057 qt

1 qt = 0.9464 L

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24
Q

Unit conversions for mass:

A

1 g = 0.002205 lb

1 lb = 453.6 g

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25
Q

1 in =

A

2.54 cm

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26
Q

Systematic error:

A

can be caused by an imperfection in the equipment being used or from mistakes the individual makes while taking the measurement. A balance incorrectly calibrated would result in a systematic error. Consistently reading the buret wrong would result in a systematic error.

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27
Q

Random errors:

A

often result from limitations in the equipment or techniques used to make a measurement. Suppose, for example, that you wanted to collect 25 mL of a solution. You could use a beaker, a graduated cylinder, or a buret. Volume measurements made with a 50-mL beaker are accurate to within 5 mL. In other words, you would be as likely to obtain 20 mL of solution (5 mL too little) as 30 mL (5 mL too much). You could decrease the amount of error by using a graduated cylinder, which is capable of measurements to within 1 mL. The error could be decreased even further by using a buret, which is capable of delivering a volume to within 1 drop, or 0.05 mL.

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28
Q

Intensive property:

A

does not depend on the size of the system or amount of material in the system (ex.; density)

29
Q

Extensive property:

A

depends on the quantity of the sample; relates to the amount of the substance present (ex.; mass and volume)

30
Q

Law of conservation of mass was proven by:

A

Antoine Lavoisier, who experimented with the thermal decomposition of mercury(II) oxide.

31
Q

The Law of Definite Proportions:

A

Also called the Law of Definite Composition–regardless of the amount, a pure compound always contains the same elements in the same proportions by mass.
Law of conservation of mass is applied to compounds–mass of the compound is equal to the masses of the elements that make up the compound.

32
Q

The Law of Multiple Proportions:

A

Also called Dalton’s Law–when one element combines with another to form more than one compound, the mass ratios of the elements in the compounds are simple whole numbers of each other.

33
Q

Compounds:

A

element + element

34
Q

Mixtures:

A
  1. Most of the matter we encounter consists of mixtures of different substances.
  2. Each substance in a mixture retains its own chemical identity and hence, its own properties.
  3. Unlike pure substances, the composition of mixtures can vary.
  4. The substances making up a mixture are called components of the mixture (ex: latte–milk, espresso, sugar).
  5. Mixtures can be separated into pure substances: elements and/or compounds (ex. air is a mixture of oxygen, nitrogen, H2O, CO2, etc.)
  6. Mixtures can be classified as homogeneous or heterogeneous.
35
Q

Homogeneous mixtures:

A

the same throughout, don’t vary in composition from one region to another.

36
Q

Heterogeneous mixtures:

A

contain regions that have different properties from those of other regions (ex. sand and water).

37
Q

Pure substances:

A

matter that has distinct properties and a composition that does not vary from sample to sample (ex. H2O and NaCl).
- All substances are either elements or compounds.

38
Q

Quantum numbers:

A

n, l, ml, ms
n = principal energy level (can not be zero)
n = 1, 2, 3, 4 …

l = sublevel (s, p, d, f)
s = 0, p = 1, d = 2, f = 3
l is less than or equal to n - 1

ml = orbital
s = 0; p = -1, 0, +1; d = -2, -1, 0, +1, +2; f = -3, -2, -1, 0, +1, +2, +3

ms = spin
+1/2 or -1/2

39
Q

Atomic emission spectra:

A

Ground state —(energy absorbed)—> excited state —(energy released)—> when energy returns to:

n = 1; emits UV rays
Called the Lyman Series

n = 2; emits visible rays
Called the Balmer Series

n = 3; emits infrared rays
Called the Paschen Series

40
Q

Auf Bau Principle:

A

each electron occupies the lowest energy orbital available

41
Q

Pauli Exclusion Principle:

A

maximum of 2 electrons per orbital

42
Q

Hund’s Rule:

A

electrons must occupy all orbitals of equal energy before pairing up.

43
Q

Properties of Group IA:

A
  1. Alkali Metals
  2. One valence electron
  3. Low ionisation energies, thus very reactive
  4. Oxidised to form +1 ion (ex. Na+)
  5. Exist in nature only as compounds
  6. Usually combine with nonmetals
44
Q

Properties of Group IIA:

A
  1. Alkali Earth Metals
  2. Two valence electrons
  3. Higher ionisation energies than Group IA
  4. Oxidised to form +2 ion (ex. Ca2+)
  5. Primarily form ionic compounds (ex. CaCO3)
  6. Are called alkali earth metals because the “Earths” of this group [lime (CaO) and magnesia (MgO)] give alkaline reactions.
45
Q

Properties of Hydrogen:

A
  1. Has 1s2 electron configuration
  2. Nonmetal that occurs as a colourless diatomic gas (H2)
  3. Ionisation energy of hydrogen is higher than Group IA so less tendency to lose electron
  4. Forms molecular compounds
  5. Also reacts with metals to form metal hydrides (H-)
46
Q

D and F block elements:

A
  1. D-block: transition metals
  2. F-block: inner transition metals
  3. A transition metal is any element whose final electron enters a d-subshell
  4. An inner transition metal is any element whose final electron enters an f-subshell
  5. period 6–Lanthanide series
  6. period 7–Actinide series
  7. Differences in properties among transition metals are based on the ability of unpaired electrons to move into the valence shell.
  8. More unpaired electrons in the d-subshell = increased hardness and increased melting point + boiling point
47
Q

Transition Metals and Inner Transition Metals:

A
  1. Across a period, little variation in atomic size, electronegativity and ionisation energy; differences in physical properties.
  2. Physical properties are determined by their electron configurations (unpaired electrons moving into the d-subshell)
48
Q

Transition Metals:

A
  1. can lose 2s electrons to become 2+
  2. can form multiple oxidation states
  3. can exhibit magnetic properties–diamagnetic [paired electrons] or paramagnetic [unpaired electrons]
  4. Fe, Co, Ni are Ferromagnetic–form permanent magnets
49
Q

Inner Transition Metals:

A
  1. Lanthanides–little variation in properties

2. Actinides–radioactive elements; only 3 exist in nature; synthetic: transuranium for atomic numbers greater than 92.

50
Q

Group IIIA–Boron family:

A
  1. do not occur elementally in nature
  2. are scarce in nature (except Al)
  3. have three valence electrons
  4. metallic (except B which is metalloid)
  5. are chemically reactive at moderate temperatures (except B)
51
Q

Group IVA–Carbon family:

A
  1. one nonmetal (C)
    - - two metals (Sn, Pb)
    - - two metalloids (Si, Ge)
  2. have 4 valence electrons, tend to form covalent bonds
  3. ability to combine with itself in long chains
  4. exists in at least 3 allotropic forms
    - -graphite and diamond
52
Q

Group VA–Nitrogen family:

A
  1. nonmetals (N, P)
    - - metalloids (As, Sb)
    - - metals (Bi)
  2. have 5 valence electrons
  3. form covalent compounds with oxidation states of +3 (ex. P) or +5 (ex. N)
53
Q

Group VIA–Oxygen family:

A
  1. has 6 valence electrons
  2. tends to form covalent compounds with other elements
  3. two molecular forms, O2 and O3
  4. O2–oxygen and O3–ozone are allotropes–different forms of the same element in the same state.
  5. tendency to attract electrons from other elements (to oxidise them)
  6. oxygen in combination with metal is almost always present as the oxide ion, O2-
  7. Sulfur also exists in several allotropic forms–S8 most common
  8. Sulfur gains electrons to form sulfides containing S2- ion.
54
Q

Group VIIA–Halogens:

A
  1. nonmetals
  2. melting and boiling points increase with increasing atomic number
  3. F and Cl are gases at room temperature
    - - Br is a liquid at room temperature
    - - I is a solid at room temperture
  4. All exist as diatomics:
    - - Halogens are very reactive
    - - They only need one electron to achieve an octet
    - - Can bond covalently with another of the same element
    - - Highly negative electron affinities (readily become anions, X-)
  5. F is most electronegative–removes electron from almost any substance
  6. Cl is most useful–it reacts slowly to form stable aqueous solutions (ex. H2O(l) + Cl2(g) –> HCl(aq) + HOCl(aq))
55
Q

Group VIIIA–Noble Gases:

A
  1. all gases and nonmetals
  2. all monatomic i.e., consist of single atoms rather than molecules
  3. completely filled s and p valence shells
  4. have the largest ionisation energies
  5. exceptionally unreactive (inert)
56
Q

Electronegativity:

A
  1. describes unequal sharing of electrons between toms
  2. it is the ability of an atom in a molecule to attract a shared electron to itself
  3. the higher the value, the more the electron in the bond resides near the atom
  4. electronegativity is measured by the polarities of the bonds between various atoms
  5. values range from 4.0 to 0.7
  6. as you move down the periodic table, the electronegativity decreases, with Cs having the lowest (0.7).
  7. as you move to the right of the periodic table the electronegativity increases, with F having the highest (4.0).
  8. Covalent bond–electrons shared equally [ex. H–H]
  9. polar covalent bond–electrons unequally shared [ex. H–Cl]
  10. ionic bond–no sharing of electrons [ex. H–F]
57
Q

Electronegativity is defined as:

A

the numerical value associated with an atom’s ability to form a covalent bond

58
Q

Electronaffinity is defined as:

A

the amount of energy that is released when the electron attaches to the atom

59
Q

Ionisation Energy:

A
  1. The minimum energy required to remove an electron from the ground state of an atom or ion
  2. has an impact on chemical behaviour
  3. The greater the ionisation energy, the more difficult it is to remove an electron
  4. Ionisation energy increases as successive electrons are removed–electrons are being pulled away from a more positive ion, which requires more energy.
  5. Ionisation energy decreases down a group, increases across a period
  6. Metals have low ionisation energies
    - - Nonmetals have high ionisation energies
60
Q

Atomic Radii:

A
  1. within each group, atomic radius tends to increase from top to bottom
  2. going down a column the outer electrons are farther from the nucleus
  3. within each period, atomic radius tends to decrease from left to right (due to Zeff – effective nuclear charge)
  4. increase in Zeff draws the valence electrons closer to nucleus, decreasing atomic radius
61
Q

Ionic Radii:

A
  1. radii of ions based on distances between ions in ionic compounds
  2. size depends on the number of electrons it possesses and on where the electrons reside
  3. formation of a cation creates ‘space’ by reducing electron-electron repulsions (cations smaller than parent atoms)
  4. the opposite is true of anions (anions larger than parent atoms)
  5. ions with same charge, size increases going down a column (Cs+ > Na+)
62
Q

Isoelectronic series:

A

Group of ions containing the same number of electrons [ex. O2-, F-, Na+, Mg2+; all have 10 electrons]

63
Q

Ionic bonds:

A

an electrostatic force that holds ions together (ratio of ions)

64
Q

Ionic compounds:

A
  1. formula units
  2. ionic crystals/crystal lattice
  3. metals (cations) & nonmetals (anions)
  4. extremely high melting point and boiling point (ex. table salt, melting point of 801 C)
  5. Lattice energy–strength of the bond (kJ)
  6. Coulomb’s Law
  7. Solubility–forms electrolytes
  8. rough in texture
65
Q

Coulomb’s Law:

A

Applied in ionic compounds.
F = K(q1q2)/r^2, where q-values represent charges.
The bigger the radius and the smaller the charges, the easier it is to break apart and melt.

66
Q

Covalent Compounds:

A
  1. molecules
  2. soft and round
  3. 2 or more nonmetals–sharing valence electrons
  4. melting point and boiling point is relatively low (ex. sugar melting point is 186 C)
  5. based on intermolecular forces
  6. solubility: doesn’t break into individual ions
67
Q

Metallic Bonding:

A
  1. two metals bonded together
    ‘sea of electrons’ – ‘delocalised electrons’
  2. electrons are shared throughout, giving the structure a sheet-like appearance with electrons surrounding the outside and protons on the inside. Electrons are being shared amongst all positive nuclei, belong to the whole metal.
  3. the melting point is reasonable, (because of fluidity with structure) ex. Sn = 232 C
  4. the boiling point is high, 2623 C, because negative charges hold the structure together extremely well.
  5. good conductors of electricity
  6. malleable–can be hammered into sheets
  7. ductile–can be made into wires
68
Q

Covalent bonds:

A
  1. strength depends on the distance between the 2 nuclei (not as strong as ionic bonds)
  2. bond length depends on number of bonds–the higher the number of bonds, the shorter the length will be and vice versa.
  3. strength of bonds are inversely proportional to length.
  4. when these bonds are created, the exothermic process takes place–releases energy, lowering the overall energy of the atoms that like to be bonded together.
69
Q

Hydrate:

A

A compound with a specific number of water molecules bound to its atom.
Examples: 1. ammonium oxalate monohydrate
2. calcium chloride dihydrate
3. sodium acetate trihydrate

Chemistry (2 decks)

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