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CHEM 119 > Unit 1: Matter > Flashcards

Unit 1: Matter Flashcards

1
Q

define matter

A

anything that has mass and volume

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2
Q

define pure substances

A

matter that is uniform throughout and cannot be separated by physical means

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3
Q

define mixture

A

two or more pure substances physically combined

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4
Q

examples of mixtures

A
  • koolaid
  • air
  • steel (alloy, solid mixture)
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5
Q

define heterogenous mixture

A

a mixture that is NOT the same throughout (ex: salad)

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6
Q

define homogenous mixture

A
  • aka solution
  • a mixture that is the same throughout (ex: koolaid)
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7
Q

define elements

A

a pure substance consisting of a single type of atom

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8
Q

define compounds

A

a pure substance that is the chemical combination of two or more elements (ex: NaCl, salt)

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9
Q

what are the two types of pure substances

A
  • elements
  • compounds
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10
Q

what is the difference between salt water and NaCl salt

A
  • salt water can be separated by physical means, is a mixture
  • NaCl salt is a pure substance and must be separated by chemical means
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11
Q

what is the simplest alcohol

A

methanol

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12
Q

is natural water containing ions and compounds a mixture or pure substance

A
  • mixture
  • homogenous (solution)
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13
Q

is distilled/deionized water a mixture or pure substance

A
  • pure substance
  • contains only water molecules and is uniform throughout
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14
Q

describe the 3 types of matter

A
  • solid: organized, rigid
  • liquid: moving, close particles
  • gas: independent particles
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15
Q

describe the volume and shape of the 3 types of matter

A
  • solid: definite volume and shape
  • liquid: definite volume, indefinite shape
  • gas: indefinite volume and shape
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16
Q

define physical properties

A

can be measured or observed without changing the composition or identity of the substance

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17
Q

examples of physical properties

A
  • density
  • conductivity
  • melting point
  • color
  • hardness
  • temperature
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18
Q

define chemical properties

A

describe the way a substance may change or react to form other substances

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19
Q

examples of chemical properties

A
  • flammability
  • corrosivity
  • reactivity
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20
Q

define chemical reactions

A

occur during chemical changes

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21
Q

what is the gas in the bubbles of pure water boiling

A
  • water (H2O)
  • boiling is a phase change (physical) so the water will stay the same substance
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22
Q

define extensive properties

A

properties that depend on the amount of matter present (scalar)

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23
Q

examples of extensive properties

A
  • mass
  • volume
  • surface area
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24
Q

define intensive properties

A

properties that do not depend on the amount of matter present

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25
examples of intensive properties
- melting point - density - temperature
26
what is the mass and charge of protons, neutrons, and electrons
- protons: 1amu, 1+ - neutrons: 1amu, 0 - electrons: 1 amu, 1-
27
what distinguishes elements from each other
number of protons
28
how do neutral atoms become ions
gaining or losing electrons
29
define cations
- positive charge - more protons - usually formed by metals
30
define anions
- negative charge - more electrons - usually formed by nonmetals
31
how was an amu previously defined
1/12 of the weight of a 12C atom
32
define isotopes
atoms with the same number of protons (atomic number) and different number of neutrons (mass number)
33
what is used to determine the existence of different isotopes and how does it work
- mass spectrometry - ionized isotopes are separated based on the difference in mass
34
what is the atomic mass
mass of the average of all the isotopes
35
how do you find atomic mass when given the weight and abundance of isotopes
- weighted average - (% abd of 1st isotope as decimal)(mass of 1st isotope) + (% abd of 2nd isotope as decimal)(mass of 2nd isotope)
36
what is the most abundant isotope in natural uranium
U-238
37
what uranium isotope is needed to use in nuclear power plants
U-235
38
what % U-235 is needed for reactor grade uranium
3-4%
39
what % U-235 is needed for weapons grade uranium
90%
40
how is uranium enrichment accomplished
- centrifuge - heavier isotopes move to the edge and the lighter isotope can be extracted
41
define periodic table
systemic catalog of elements arranged in order of atomic number
42
what are the rows and columns of a periodic table called
- rows: periods - columns: groups
43
what is the name of group 1 in the periodic table
alkali metals
44
what is the name of group 2 in the periodic table
alkaline earth metals
45
what is the name of group 6 in the periodic table
chalcogens
46
what is the name of group 7 in the periodic table
halogens
47
what is the name of group 8 in the periodic table
noble gases
48
define chemical compounds
formed from fixed ratios of atoms or ions
49
define ionic compounds
- usually formed between cations (metals) and anions (nonmetals) - can also be formed between polyatomic ions
50
define covalent
- formed between nonmental atoms
51
define empirical formula
- simplest whole number ratio of atoms of each element in a compound - always used for ionic compounds
52
define molecular formula
- exact number of atoms of each element in an individual molecule - covalent compounds have molecular formulas
53
define structural formula
drawing of the connectivity of the atoms
54
example of the empirical and molecular formula for hydrogen peroxide
- empirical: HO - molecular: H2O2
55
how are ionic compounds formed
metal transfers electrons to a nonmetal
56
what is the charge of elements in groups 1, 2, 6, 7
- 1: 1+ - 2: 2+ - 6: 2- - 7: 1-
57
what is the charge of Al, Zn, Cd, Ag
- Al: 3+ - Zn: 2+ - Cd: 2+ - Ag: 1+
58
describe the naming system of oxyanions in the group halogens (number of oxygens and prefixes)
- know the number of oxygens in the -ate form - one less oxygen in the -ite form - two less oxygens in the hypo-ite form - one more oxygen in the per-ate form - hypo-ite to ite to ate to per-ate
59
what order do you write ionic compounds in
- metal cation first - nonmetal anion second
60
describe how you write the formula for ionic compounds
- metal cation first and nonmetal anion second - make the charges equal by adding subscripts (ex: +2=-2) - make sure subscripts are in empirical formula
61
describe how to name an ionic compound
- write metal cation first - write nonmetal anion: if element change ending to -ide, if polyatomic ion write its name
62
what must you include in the name of a compound where the cation can have more than one possible charge
- write roman numeral of charge in parenthesis - ex: iron(III) nitrate
63
define molecule
aggregate of at least two atoms held together by covalent bonds
64
define covalent bonds
electrons are shared between nonmetal atoms
65
what formula should you write covalent compounds in
molecular or structural formulas
66
define diatomic elements and list them
- elements that naturally occur as molecules containing two atoms (O with subscript 2) - H2, N2, O2, F2, Cl2, Br2, I2
67
define acid
- covalent - contains H+ and an anion - give off H+ in water
68
define inorganic compounds
- covalent - contains nonmetals
69
define organic compounds
- covalent - contains C and H (sometimes O)
70
how do you know the amount of H+ to have in an acid
H+ must balance with the charge of the anion
71
acids must have balanced charges so does this mean they are ionic compounds
- no - acids are covalent not ionic
72
explain how to name acids
- anion ends with -ide: hydro_ic acid (HF, hydrofluoric acid) - anion ends with -ate: _ic acid (HNO3, nitric acid) - anion ends with -ite: _ous acid (HNO2, nitrous acid)
73
explain how to name binary covalent compounds
- prefix used to denote number of atoms of each element in the compound - don't use mono- for first element
74
what 3 things does the mass number of an element represent
- average weight of a single atom in amu - average weight of one mole of atoms in grams - average weight of 6.022E23 atoms in grams
75
define avogadro's number
- 6.02E23 - number of atoms in 1 mole
76
how many moles of oxygen are in 1 mole of glucose, C6H12O6
6 mol O
77
how do you find the molar mass of a molecule
multiply the number of atoms of each element by its atomic mass and add them all together
78
what measurement is needed to figure out how many moles of glucose are in 10g
- molar mass of glucose - 180.12g = 1 mole glucose
79
define solution
- homogenous mixture of two or more substances - can be liquid, solid, or gaseous
80
define solvent
substance present in the greatest quantity in a solution
81
define solute
- substances other than the solvent in a solution - solutes are dissolved in the solvent
82
define aqueous solution
liquid solution in which the solvent is water
83
define concentration
amount of solute present in a given quantity of solvent or solution
84
define molarity
- molar concentration moles of solute per liter of solution
85
explain solution preparation
- measure the amount of solute you want - funnel solute into volumetric flask - add a portion of the solvent to the flask - swirl the mixture until all of the solute is dissolved - add additional solvent needed
86
why shouldn't you pour all of the solvent into the solute when creating a solution
- the solute will take up volume as it is dissolved in the solvent - for precise measurements of a solution you must wait for the solute to dissolve and take up volume before adding the rest of the solvent
87
define dilution
procedure for preparing a less concentrated solution from a more concentrated (stock) solution
88
what is the equation for dilution
- M1V1 = M2V2 - molarity of concentrated times volume of concentrated = molarity of diluted time volume of diluted

CHEM 119 (8 decks)

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