Chapter 1-4, 6 Flashcards
(119 cards)
1
Q
Physical Change
A
A change that alters only the state or appearance of a substance, but not its chemical composition
Ex: change in state from liquid to gas
2
Q
Chemical Change
A
A change that alters the molecular composition of a substance
3
Q
Physical Property
A
A property that a substance displays without changing its chemical composition
Ex: melting point
4
Q
Chemical Property
A
A property that a substance displays only by changing its composition via a chemical change
Ex: combustibility
5
Q
Kinetic Energy
A
Energy associated with motion
6
Q
Potential Energy
A
Energy associated with its position or composition
7
Q
Thermal Energy
A
Energy associated with the temperature of an object
8
Q
Law of Conservation of Energy
A
Energy can neither be created nor destroyed, only converted from one form to another
9
Q
Units
A
The standard quantities used to specify measurements
10
Q
International System of Units (SI)
A
The standard unit system used by scientists, based on the metric system
11
Q
How do you measure something?
A
Measure to the number of markings on the measurement device then estimate to the next digit
12
Q
Significant Figures
A
In any reported measurement, the place holding digits that indicate the precision of the measured quantity
* the greater the number of sig figs, the greater the certainty of the measurements
13
Q
Name the sig fig rules
A
- all non zero digits are significant: 28.03 has 4 SF
- zeros between non zero digits are significant: 408.1 has 4 SF
- zeros to the left of the first non zero digit are not significant (they only serve to locate the decimal point): 0.0032 has 2 SF
- Zeros after a decimal point are always significant: 45.000 has 5 SF
- zeros before a decimal point and after a non zero digit are always significant: 140.00 has 5 SF
- Zeros before an implied decimal point are ambiguous and scientific notation should be used
14
Q
Exact numbers
A
Numbers that have no uncertainty and thus do not limit the number of SF in any calculation
Ex: 3 atoms means 3.0000000… atoms
15
Q
SF in multiplication and division calculations
A
the result carries the same number of SF as the factor with the fewest SF
16
Q
SF in subtraction and addition calculations
A
the result carries the same number of decimal places as the quantity with the fewest SF
17
Q
SF in logarithmic and natural calculations
A
the result has the same number of SF after the decimal point as there are SF in the number whose log is being calculated. The digits before the decimal place in the answer represent the order of magnitude
Ex: log(4.5250 x 10^1)= 1.65562
the same is true for natural calculations
18
Q
SF in antilogarithm calculations
A
the number of SF in the result is the same as the number of SF after the decimal place. The number before the decimal place represents the order of magnitude.
Ex: 10^6.125=1.33x10^6
19
Q
Precision
A
refers to how close a series of measurements are to on another or how reproducible they are
20
Q
Accuracy
A
refers to how close the measure value is to the actual value
21
Q
Law of Conservation of Mass
A
matter can neither be created no destroyed, only rearranged
- atoms combine in simple whole number ratios to form compounds
22
Q
Atomic Theory (John Dalton)
A
- Each element is composed of tiny indestructible particles called atoms
- all atoms of a given element have the same mass and physical/ chemical properties that distinguish them from the atoms of other elements
- atoms combine in simple whole number ratios to form compounds
- atoms of one element cannot change into atoms of another element. In a chemical reaction, atoms only change the way they are bound together with other atoms
23
Q
Electron
A
a negatively charged, low mass particle found outside the nucleus
24
Q
Nucleus
A
the very small, dense core of the atom that contains most of the atom’s mass and all of its positive charge, contains protons and neutrons
25
Proton
positively charged subatomic particle found in the nucleus
26
Neutron
a electrically charged neutral subatomic particle found in the nucleus, with a mass almost equal to that of a proton
27
Atomic Number (Z)
the number of protons in a nucleus
28
Chemical Symbol
a unique one or two letter abbreviation for each element
29
Isotope
atoms of the same element with the same number of protons, but different number of neutrons and different masses
30
Natural Abundance
the relative percentage of a particular isotope in a naturally occurring sample with respect to other isotopes of the same element
31
Mass Number (A)
the sum of the number of neutrons and protons in an atom
32
Atomic Mass
the average mass in u of the atoms of a particular element based on the relative abundance of the various isotopes; it is numerically equivalent to the mass in grams of one mole of the element
33
Nuclear theory
the theory that most of the atom's mass and all of its positive charge is contained in a small dense nucleus
34
Unified atomic mass unit (u)
a unit used to express the masses of atoms and subatomic particles
35
ion
an atom or molecule with a net charge cause by the loss of gain of electrons
36
cation
a positively charged ion
37
anion
a negatively charged ion
38
mole
a unit defined as the amount of material containing 6.022 x 10^23 (avogadro's number) particles
39
Molar mass
the mass (g) of one mol of atoms of an element usually equivalent to the atomic mass of the element in u
molar mass= 1 mol
molar mass/ 1 mol
40
Converting between mols and atoms
1 mol/ 6.022 x 10^23
or
6.022 x 10^23/1 mol
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main group metal
tends to lose electrons and forms cations with the same number of electrons as the nearest noble gas
42
Main group non metal
tends to gain electrons and form anions with the same number of electrons as the nearest noble gas
43
ionic bond
- occur between metals and non metals
| - involves the transfer of electrons from one atom to another
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covalent bond
- occur between 2 or more non metals
| - involves the sharing of electrons between 2 atoms
45
Chemical formula
a symbolic representation of a compound which indicate the elements present in the compound and the relative number of atoms of each
46
empirical formula
a chemical formula that shows the simplest whole number ratio of atoms in the compound
47
molecular formula
a chemical formula that shows the actual number of atoms of each element in a molecule of a compound
48
structural formula
a molecular formula that shows how the atoms in a molecule are connected or bonded to each other
49
Polyatomic ion
an ions composed of two or more atoms
50
Ion compound
- ionic compounds always contain positive and negative ions
- in a chemical formula, the sum of the charges of the cations must equal the sum of the charges of the anions
- a formula reflects the smallest whole number ratio
51
How do your write the formula for ionic compounds
- write the symbol for the metal cation and its charge followed by the symbol for the non metal anion and its charge
- adjust the subscript on each cation and anion to balance
- check the sum of the charges
Al^3+. O^2-
Al2O3
Cations- 2(3+)= 6+
Anions- 3 (2-)=6-
52
How do you name ionic compounds?
-identify the compound as an ionic compound (metal +non metal)
- name the metal first, if the metal can form cations with different charges, add the charge in roman numerals
- name the non metal and add the suffix "ide"
CrBr3
Chromium (III)
Chromium (III) bromide
53
How do you name polyatomic ionic compounds?
- identify the compound as a polyatomic ionic compound (metal +polyatomic0
- use the same rules as naming normal ionic metals
- use the name of the polyatomic ion
NaNO2
Sodium
Sodium nitrate
* if the compound contains both a polyatomic cation and anion use the names of both
NH4NO3
ammonium nitrate
54
Acid
a molecular compound that donates H+ ions when dissolved. Composed of hydrogen and one or more non metals
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binary acids
composed of hydrogen and a non metal
56
how do you name binary acids
hydro + base name of the non metal + "ic acid"
hydrobromic acid
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oxyacid
composed of hydrogen and an oxyanion (non metal + oxygen)
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how do you name an oxyacid
oxyanion ending in "ate'- base name of axyanion + "ic acid"
- nitric acid
oxyanion ending in "ite' - base name of oxyanion + "ous acid"
- sulfurous acid
59
Formula Mass
the average mass of a molecule
60
Mass Percent Composition
an element's percentage of of the total mass of a compound containing an element
mass percent of element x = (mass of element x in 1 mol of compound/ mass of 1 mol of compound)(100)
61
Empirical Formula Molar Mass
the sum of the masses of all atoms in an empirical formula
62
Combustion Analysis
a method of obtaining empirical formulas for unknown compounds, especially those containing carbon and hydrogen, by burning a sample of the compound in pure oxygen and analyzing products of the combustion reaction
63
Chemical reaction
a process in which one or more substances are converted into on or more different ones
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Combustion reaction
a type of chemical reactions in which a substance combines with oxygen to form one or more oxygen containing compounds
65
chemical equation
a one line expression that shows the reactants and products in a chemical reaction
66
reactant
substances that are consumed in a chemical reaction
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products
the substances produced in a chemical reaction; they appear on the right hand side of the chemical equation
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solvent
the majority component of a solution
69
solution
a homogenous mixture of 2 substances
70
solute
the minority component of a solution
71
aqueous solution
a solution in which water acts as the solvent
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molecular equation
an equation showing the complete neutral formula for each compound in a solution
73
complete ionic equation
an equation which lists individually all of the ions present as either reactants or products in a chemical equation
74
spectator ions
ions in a complete ionic equation that do not participate in the reaction and therefore remain in the solution
75
net ionic equation
an equation that shows only the species that actually change during the reaction
76
acid base reaction
a reaction in which an acid reacts with a base and the 2 neutralize each other, producing water
77
neutralization
the state of a solution where the concentrations of H3)+ and OH- are equal
78
Arrhenius definitions of acids and bases
the definitions of an acid as a substance that produces H+ ions in aqueous solution and a base as a substance that produces OH_ ions in an aqueous solution
79
hyrdronium ion
H3O+ the ion formed from the association of a water molecule with an H+ ion donated by an acid
80
polypro tic acid
an acid that contains more than on ionized proton and releases them sequentially
81
oxidation reduction reaction
reactions in which electrons are transferred from one reactant to another and the oxidation states of certain atoms are changed
82
oxidation
the loss of one or more electrons; also the gaining of oxygen or the loss of hydrogen
an increase in oxidation state
83
reduction
the gaining of one or more electrons; also the gaining of hydrogen or the loss of oxygen
a decrease in oxidation state
84
oxidation state
a positive or negative whole number that represents the charge an atom in a compound would have if all shared electrons were assigned to the atom with a greater attraction for those electrons
85
oxidation agent
oxidizes another substance (is reduced)
86
reducing agent
reduces another substance (is oxidized)
87
limiting reactant
the reactant with the smallest stoichiometric amount in a reactant mixture and consequently limits the amount of product in a chemical reaction
88
theoretical yield
the greatest amount of product that can be made in a chemical reaction based on the limiting reactant
89
percent yield
the percentage of the theoretical yield of a chemical reaction that is actually produced
90
actual yield
the amount of product actually produced by a chemical reaction
91
molarity (M)
a means of expressing solution concentration as the number of moles of solute per litre of solution
92
Stock solution
a highly concentrated form of a solution used in laboratories to make less concentrated solutions via dilution
93
concentrated solution
a solution that contains a large amount of solute relative to the amount of solute
94
dilute solution
a solution that contains a very small amount of solute relative to the amount of solvent
95
thermochemistry
the study of the relationship between chemistry and energy
96
energy
the capacity to do work
97
heat (q)
the flow of energy caused by a temperature difference
98
thermal energy
a type of kinetic energy associated with the temperature of an object, arising from the motion of individual atoms or molecules
99
kinetic energy
energy associate with motion of an object
100
potential energy
energy associate with the position or composition of an object
101
chemical energy
energy associated with the relative positions of electrons and nuclei in atoms and molecules
102
law of conservation of energy
energy can neither be created or destroyed, bu it can be transferred from one object to another
103
system
a specific area of interest
104
surroundings
everything else
105
the first law of thermodynamics
the total energy in the universe is constant
106
internal energy (u)
the sum of kinetic energy and potential energies of all of the particles that compose a system
107
state function
a function whose value depends only on the state of the system, not on how the system got to that state
108
Heat capacity (c)
the quantity of heat required to change a system's temp by 1 degree
109
specific heat capacity
the amount of heat required to raise the temp of 1 gram of a substance by 1 degree
110
pressure volume work
work that occurs when a volume change takes place against an external pressure
111
bomb calorimeter
a piece of equipment designed to measure delta u for combustion reactions at constant volume
112
enthalpy (H)
the sum of the internal energy of a system and the product of its pressure and volume; the energy associated with the breaking and forming of bonds in a chemical reaction
113
exothermic reaction
a chemical reaction that releases heat to its surrounds
delta h is less than 0
114
endothermic reaction
a chemical reaction that absorbs heat from its surroundings
delta h is greater than 0
115
coffee cup calorimeter
a piece of equipment designed to measure delta h for reactions at a constant pressure
116
hess's law
the law stating that if a chemical equation can be expressed as the sum of a series of steps, then delta h for the overall equation is the sum of the enthalpies of reactions for each step
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standard state
for a gas the standard state is the pure gas at a pressure of exactly 1 bar; for a liquid or solid the standard state is the pure substance in its most stable form at a pressure of 1 bar and the temperature of interest; for a substance in solution the standard state is a concentration of exactly 1 mol/L
118
standard enthalpy change
the change in enthalpy for a process when all reactants and products are in their standard states
119
standard enthalpy of formation
the change in enthalpy when 1 mol of a compound forms from its constituent elements in their standard states
