Midterm Flashcards
(135 cards)
1
Q
What is chemistry
A
Chemistry is the science of the properties structure and transformation of matter.
2
Q
What is matter
A
Matter is anything with mass and takes up space
Matter cannot be destroyed
3
Q
What’s the law of conservation of mass.
A
Lavoisier (1743-1794)- is a chemical reaction, matter is neither created nor destroyed it is transformed into something else
4
Q
Chemical change/ chemical reaction
A
Substances are used up and other compounds are formed to take their place
5
Q
Physical change
A
Identity of matter remains the same but involves change in its state
6
Q
States of solids
A
They have a definite shape and a definite volume
All partials are very close together
7
Q
States of liquids
A
Liquids have an indefinite shape and a define volume
Partials are not very close together
8
Q
States of gasses
A
Gasses have an indefinite shape and an indefinite shape
Partials and very far apart and move freely
9
Q
Why study the state of matter of a substance
A
To understand how things form/exist
Everything is matter
It’s the building blocks to everything
Risk assessment
Transportation
Handling
Reaction potential
10
Q
Physical property states
A
Solid
Liquid
Gas
Plasma
11
Q
Physical property change of state
A
Melting
Freezing
Boiling/vaporizing
Condensation
Ionization
Relaxation
12
Q
Physical property characteristics
A
Colour
Shape
Size
Texture
Mass
Volume
Density
13
Q
Chemical property signs of change
A
Order
Change in temp
Change in colour
Bubbles form
Solids form
14
Q
Chemical property chemical reactions
A
Atoms from two substances combine to form new substances
15
Q
Chemical property examples
A
Burning
Rusting
Rotting
Tarnishing
16
Q
What does EC stand for
A
Electrical conductivity
17
Q
What does DO stand for
A
Dissolved oxygen
18
Q
Pure substance-element
A
A pure chemical substance consisting of one type of atom which can’t be broken down further
118 elements
19
Q
Pure substance-Compound
A
Define and constant composition same properties under a certain set of conditions
Water (H2O) 11.2% by mass H and 88.8% by mass O2
20
Q
Homogeneous mixture
A
Uniformed composition and properties though the sample
Difficult to see impurities in homogeneous mixture
Ex. Air, gas, salt solution
21
Q
Heterogeneous mixture
A
Non-uniform composition and properties throughout the sample. Different components are visibly distinguishable from one another
Ex. Water and oil
22
Q
How do you test a theory
A
Facts-statement based on direct experiences
Hypothesis-proposed statement to explain facts but lacks proof
Theory- hypothesis with some degree of proof. Establishes a cause and effect relationship
23
Q
How many significant figures is this 4.1658 and 8.45
A
4.1685- 5 significant figures
8.45- 3 significant figures
24
Q
Are zeros significant before the decimal
A
No
Ex. 0.004066
25
What BEDMAS
Bracket, Exponents, Division, Multiplication, Addition and Subtraction
26
Addition and subtraction rule
- the final value must have only as many decimal places as the least precise measurement with the least number of decimal places
- the answer cannot have more digits to the right of the decimal points that any of the original numbers
27
Multiply and divide rule
-The final value can only have as many significant figures as the least precise measurement with the least number of significant figures
-the number of significant figures in the products or quotient is the determined by the least precise measurement that have the fewest significant figures
28
What is a measurement
A method of determining a physical quantity such as length, time and temperature
29
Why do we need measurements
So we quantify time, distance, plan, schedule, risk, assessment, impacts make decisions
30
Quantity of measurements -accuracy
-how close a measurement is to the true value
-describes the difference between the measurement and the parts actual value
31
Quantity of measurement- precision
-how reproducible is the measurement
-describes the variation you see when you measure the same part repeatedly with the same device
32
What three measurement systems are used
Metric-world wide- meter, litre, gr, second, Celsius and mol
English system- United States- pounds, inches, gallons and yards
International system of units (SI)- scientific organizations- based on metric system
33
What unit is comparable in all three measurement types
Seconds
34
What systems use the power of tens for large or very small numbers
Metric system and SI
35
Factor label method
Set up equation using given units and conversion factor so that the unwanted units cancel remaining desired units
Given x conversion factor=desired
36
Does chemical identity change when substance converts states
No
37
What’s the density equation
D=m/v
38
Define density
Physical property constant at a given temperature
39
Is liquid water denser than solid water
Yes- ice represents an open crystal structure of hydrogen bonds. Which is lighter that water due to physical structure.
40
Specific gravity equation
Density of substance(g/mL)/ density of water (g/mL)
Specific gravity is unit less
41
Application of specific gravity of chemical in the environment
Specific gravity>1—> lighter that water= floats
42
Define energy
Capacity to do work
2 forms= kinetic and potential
43
What is kinetic energy
Energy motion
Possessed by any moving object
Light, heat, mechanical energy
KE=1/2 m. V2
44
What is potential energy
Stored energy
Capacity to move or cause motion due to position
Chemical and nuclear energy
45
Law of conservation of energy
Energy can neither be created nor destroyed. It is converted
46
Equation for specific heat
Q=mc🔺t
47
Democritus atom theory
Mater is made of very small indestructible units called atoms
He used the word atomos which means not to cut(Greek)
Was the first philosopher to use the word indivisible
48
Define elements
A substance that consists of the same kind of atoms
49
What are the 6 essential elements of living organisms
Carbon
Hydrogen
Nitrogen
Oxygen
Phosphorus
Surfer
50
Define compound
A substance with fixed ratio of elements
Chemical properties differ from elements
20 million known compounds
Characterized by its formula
51
Compound formulas-Combining Ratios
Formula indicates the atomic symbol of each element and subscript indicates the ratio of an element
HCl
52
First postulate of Daltons atomic theory
All matter is made up of very tiny indivisible partials called atoms
53
Second postulate of Daltons atomic theory
All atoms of a given element have the same chemical properties and atoms of different elements have different chemical properties
54
Third postulate of Daltons atomic theory
In ordinary chemistry reaction No atom of any element can disappear or change into an atom of a different element
55
Fourth postulate of Daltons atomic theory
Compounds are formed by the chemical combination of two or more different kinds of atoms. In a given compound the relative number of each kind of element are constant
56
Fifth postulate of Daltons atomic theory
A molecule is a tightly bound combination of two or more atoms that acts as a single unit
57
Law of conservation of mass
Total mass of a matter at the beginning and end of an ordinary chemical reaction is the same
Matter can neither be created nor destroyed
58
Law of constant composition
The theory Joseph Proust made that support Daltons claims
Any compound is always made up of elements in portion by mass
59
Monoatomic elements
Consist of single atoms that are not connected to each other
60
Diatomic elements
Two atoms in each molecule connected by a chemical bond O2 ⭕️⭕️
Under normal conditions free atoms don’t exist for these elements
61
Polyatomic element
More than two atoms in each molecule connected by chemical bonds
Ex O3 ⭕️⭕️⭕️⭕️
62
What does a atom consist of
Protons
Neutrons
Electrons
63
Where are protons and neutrons
In the nucleus
64
Where are electrons
On the valence rings
65
Define mass number
Mass number is one way to describe an atom
It is used to express the relative masses of elements in a compound
66
Define atomic number
Number of protons in the nucleus
All atoms of the same element have the same number of protons
118 elements=118 atomic numbers
Neutral atoms: number of electrons= number of protons
67
What are isotopes
Isotopes are an element that have the same number of protons but different number of neutrons
Isotopes of an element have almost identical properties except radioactive properties
Protons #= electrons#
68
Do isotopes have the same Z number (atomic number) as # of A (atomic mass)
No. Isotopes have the same z but different A
69
Where are isotopes found
Most elements are found on earth as mixtures of isotopes most relatively constant ratio
70
What is atomic mass number
The mass of a specific isotope
Total mass p+n in an isotope
Measured in amu
= the weighted average of masses of all isotopes
This is given on the period table
71
Predicting relative isotopes abundance
Which of the two apropos of antimony is the most abundant in nature: sb-121 (atomic mass 120.9amu) or sb-123 (atomic mass 122.9 amu)
Step 1: look up atomic weight of sb on the periodic table
=121.760 amu
Step 2- compare atomic weight to the isotopes atomic mass see which is closer to the value
Sb-121 120.9amu 121.760-120.9=0.86
Sb-123 122.9amu 122.9-121.760=1.40
Sb in nature is 121.760 amu
Sb-121 is closest to nature
Sb121 =(sb-sb121/ isotopic mass 121-isotopic mass123)x100
72
Calculating % by weight in compounds
Step 1-add the molar mass of each element (i) to find the compound molecular weight
Step 2-divid the mass of the element in the compound by the compound’s molecular weight. Multiply by 100% to get percentages
73
Percent abundance of a isotope equation
% abundance of isotope=(atomic weight of an element- atomic mass of isotope 2/atomic mass of isotope 1- atomic mass of isotope 2) x100
74
How is a periodic table organized
Families or groups=vertical rows
Periods= horizontal rows
75
Main groups of elements
1A or iupac 1
2A or iupac 2
3A-8A or iupac 13-18
Transition elements 3B-12B
Inner transition elements 58-71 and 90-103
76
Define metals
All but 24 elements are metals
Shiny and ductile
Solids at room temperature
Conduct electricity
Tend to give up electrons
Positive charged
77
Define non metals
18 metals
Do not conduct electricity
Solid, liquid or gas at room temperatures
CHNOPS-organic and biochem
Tend to accept electrons
78
Define Metalliods
6 elements
Metal and non metal properties
79
Periodicity in the period table
Properties vary in regular ways as you move up or down a column
Halogens boiling points increase as you move down
Alkali metals softness increases as you move down
80
Mass and size of an atom
Mass of an atom is concentrated in the nucleus (neutrons and electrons are in the nucleus)
Size of an atom is dictated by the electron cloud
81
How to calculate electrons in neutral atoms
Number of electrons=number of protons
82
Niels Bohr theory
Electron energy is quantized
Ground state=lowest energy level
83
Electron distribution
Electrons don’t move freely around the nucleus
Confined to specific regions=principles, energy, levels or shells
Shells are numbered from inside out
84
How many electrons can each leave hold
1st level-2 electrons
2nd level-8 electrons
3rd level- 18 electrons
4th level-32 electrons
5th level-50 electrons
85
Possible changes of potential energy in electron distribution
An electron can move from one level to another only if the energy gains or loses is exactly equal to the difference in energy between the two levels
86
Describe sub shells
Shells are divided into sub shells (s,p,d,f)
Within each shell electrons are grouped into orbital experiments
Each orbital can hold maximum of 2 electrons
87
Define s orbital shapes
Sphere shape
88
Define p orbitals shape
Dumbbell shape and 90 degrees apart from the centre axis (Px, Py, Pz)
89
Electron configuration of atoms
It’s the description of the orbitals that it’s electron occupy
Orbitals available to all atoms are the same regardless of elements
-1s, 2s, 2p, 3s, 3p, ect. ( how they increase)
90
Rule 1 for the ground state configuration
Rule 1: orbitals fill in the order of increasing energy from lowest to highest
91
Rule 2 for ground state configuration
Each orbital can hold up to two electrons with opposite spins with one arrow pointing up and the other pointing down
92
Rule 3 for ground state configuration
Hudes rule: when there is a set of orbitals of equal energy, each orbital becomes half-filled before any are completely filled
93
Showing electron configuration noble gasses
Noble gas notation-abbreviated way to show electron configuration of an atom
Uses “previous” noble has to represent all the configuration up to that gas and then show all the subsequent electrons
94
Example of expanded electron configuration
Carbon: 1s2,2s2,2px1,2py1
95
Example of condensed electron configuration
Carbon 1s2,2s2,2p2
96
Example of noble gas notation electron configuration
Carbon (He)2s1,2p2
97
Valence electrons
Electrons in the outer most shell of an atom
98
Lewis Dot structure
A way represent the valence electrons of an atom with dots around the chemical symbol
Number of dots=number of valence electrons
99
Atomic size and radius
Atomic size is determined by the size of its outer most occupied orbit
Atomic radius= atomic mass/ 2
100
Atomic radius-periodic trend
Atomic radii (ionic radius) of elements increases Down the groups (columns)
Atomic radii (ionic radius) of elements increases to left across the period (rows)
101
What happens has you go down a group
Ionic radius increases
Number of shells increase
More shells of electrons
Shielding increases
Also less attraction means electrons are lulled in less by the nucleus
Atomic radius increases
First ionization energy decreases
102
Ions cations and anions
Atoms can gain or lose electrons
Ion= atom with unequal number of protons and electrons
Anion= ion with negative charge
Cation= ion with positive charge
103
Types of cations
Formed when electrons are removed from an atom
Positive ions are always smaller then the neutral atom
104
Types of anions
Forms when an atom gains electrons
Negative ions are always larger than the neutral atom
105
Ionization energy
Ionization energy-a physical property of elements
It is the energy required to remove the most loosely held electron from an atom in the gaseous phase
The more difficult to remove an electron the higher the ionization energy
Ionization energy is always positive because energy must be supplied to remove the attraction force between the nucleus and electrons
106
Ionization energy and periodicity
IE increases as you go further up the column (bottom to top)
- electrons further from nucleus are shielded by inner electrons and less attracted to the nucleus therefore less IE is required to remove the outer electrons
IE increases as you go across a row ( left to right)
-valence electrons across a row are in the same shell but number of protons in nucleus increases outer electrons attraction to nucleus and makes them harder to remove
107
Octet rule
Lewis observed that the noble gasses lacked chemical reactivity
Thus lack of reactivity indicated a high degree of stability in the noble gas electron configuration
The tendency of atoms to react in different ways to achieve an outer shell configuration of 8 electrons like noble gasses became known as octet rule
Most common among 1A-7A
108
Octet rule ions
Atoms with close to 8 valence electrons tend to gain electrons
Gain electrons=negative charge=anions
Atom with 1 or 2 valence electrons tend to lose electrons
Loose electrons=positive charge=cation
Comparing an ion to the original atom #p, and #n are the same only the #e charges in the valence shell
109
First expectation of octet rule
Ions of period 1 and 2 with charge greater that +2 are unstable and therefore ions of there elements don’t exist in nature
110
2nd exception to octet rule
Octet rule does not apply to group 1B-7B elements (transition elements)
Most of 1B-7B from ions of 2 or more positive charges
111
Properties of atoms and their ions
Atoms and their ions have different properties
When Na and Cl are ionized their reactivity decrees and they are stable
When two atoms mix to create an ion it changes their reactivity
112
Naming cations rule 1
Elements of groups 1A, 2A and 3A form only one type of cation
The name of the cation is the name of the metal followed by the ions
Ex. Hydrogen ion
113
Naming cations rule 2
Two types of cations formed:
Systematic names: use a Roman numeral enclosed in parentheses following the name of the element to show the charge
Older common system: use the suffix-ous to show the smaller positive charge and the suffix -Ic to show larger positive charge
114
Naming anions
For monatomic (containing only one atom) anions add ide to the stem name
Ex. Bromide
115
Chemical bonds
Atoms born together so that each atom participating in the bond acquires a valence shell electron configuration the same as that of noble gas nearest to it in atomic number
116
Ionic bonds
Results from the force of attraction between cations and anions
117
Covalent bonds
Results from the force of attraction between 2 atoms sharing one or more pairs of electrons
118
How compounds are formed ( daltons 4th postulate)
Compounds are formed when atoms of more that one element combine a given compound always the same relative number and kind of atom
Compounds are tightly bonded groups of atoms held together by forces of attraction called chemical bonds
119
Ionic bond formation
Forces of attraction between a cation and anion
One atom donates electors another atom
Generally between metals and non metals
High boiling and melting points
Requires a lot of energy to break ionic bond
Electronegativity difference >1.9
120
Define electronegativity
A measure of an atoms attraction for the electrons it shares in a chemical bond with another atoms
121
Ionic compounds
Name for the entity formed by electrostatic attraction of positive and negative ions
Elements with higher EN gains electrons
Elements with lower EN donates electrons
122
Rule 1 binary ionic compounds
When only one oxidation state: cation name+ stem of anion with the suffix ide
123
Define binary
Contains two elements
124
Rule 2 naming binary ionic compounds
When >1 oxidation state:cation name+Roman numeral for charge+ stem of anion name with the suffix ide
125
Common naming rule binary ionic compounds
When >1 oxidation state:cation Latin root+ous or Ic suffix+ stem of anion name with the suffix ide
126
Naming polyatomic ionic compounds
Naming polyatomic ion as a unit
Cation name + anion name
127
Predicting ionic compound formula
Strategy: sum of positive charge=sum of negative charge
To product formula simply balance # atoms so cation and anions are equal
Subscripts represent ratios reduced to lowest whole number
128
Ionic compound formula groups and their charges
Group 1 ions ( alkali metals) have +1 charges
Group 2 ions (alkaline earth metals) have +2 charges
Group 6 ions (non metals) have -2 charges
Group 7 ions (halogens) have -1 charges
129
Covalent bond formation
Force of attraction between 2 atoms sharing one or more pairs of electrons to attain valence electrons configuration of noble gas
130
What are covalent bonds are generally between what?
Non metals and non metals
Non metals and metalloids
Lower boiling and melting points require less energy to break down a covalent bond electronegativity less than 1.9
131
Polar covalent bonds
EN difference=0.5 to less than 1.9
Unequal sharing of electrons partial charges on atoms
132
Non polar covalent bonds
EN difference=less than 0.5
Atoms share electrons equally
133
Electronegativity of noble gasses
Noble gasses have No electronegativity
Research showed that relatively stable compounds and both Xe and Kr exist in nature with F, Cl and or O
134
Sigma bonds
The orbitals involves in the bond face each other (stronger bond) (single bond/saturated)
135
Pie bonds
The orbitals involve in the bonds are parallel to each other ( not as sting as sigma) ( single or triple bonds/ unsaturated)
