Unit B (Matter, periodic table, chemical bonding, nomenclature) Flashcards
(98 cards)
1
Q
Extensive properties def’
A
A physical property that is dependent on the amount/size of substance
2
Q
Intensive
A
A physical property that’s independent of size/amount of sub
3
Q
Endothermic reaction
A
When the system (substance) absorbed heat from the surrounding, leaving the surrounding surface feeling cold.
Ex . Ammonium nitrate
4
Q
Exothermic reaction
A
the system release heat into the surroundings, making the surrounding surface feel hot
5
Q
Is photosynthesis, cooking an egg, combustion, rain an exothermic or endothermic reaction?
A
Endo, Endo, exo, exo.
Rain: Condensation of water vapor into rain releasing energy in the form of heat is an example of an exothermic process.
6
Q
What’s the mass of a proton?
A
1.0073 u (atomic mass unit)
7
Q
What’s the mass of a neutron?
A
1.0087 u
8
Q
what’s the mass of electron?
A
5.84 x 10^-4 u
9
Q
where is the electron located?
A
in the orbitals/ electron cloud
10
Q
electrons, neutrons, and protons are examples of what?
A
subatomic particles
11
Q
what’s the size of protons, electron, and neutron?
A
proton = neutron fo 1x10^-15 fm (femtometer), electrons are 1x10^-18 fm (super small)
12
Q
what’s the relationship between how the period table is organised with electrons?
A
the periodic table is organised based on how electrons are occupying the energy levels
13
Q
What’s the charge of a neutron?
A
neutral (none)
14
Q
What’s the non-metal that’s a liquid at SATP?
A
bromine
15
Q
what’s the metal liquid at room SATP?
A
mercury
16
Q
What are physical changes?
A
Changes of state, temperature, volume, dying colour, dissolving
17
Q
What are chemical changes?
A
oxidation, combustion, decomposition, new sub. formed, energy change, temp. change, precipitate
18
Q
Examples & def’of physical properties
A
A physical property is a characteristic that can be observed or measured without changing the substance’s identity.
solubility, mass, density, boiling point, state of matter, conductivity, luster
19
Q
Examples & def’of chemical properties
A
A chemical property describes a substance’s ability to change into a new substance through a chemical reaction.
reactivity, flammability, toxicty, acidity/basicity (pH), corrosiveness
20
Q
_____ and _____ independently came to the conclusion about how elements a=should be grouped.
A
Dmitri Mendeleev , Lothar Meyer
21
Q
Mendeleev predicted the discovery of ____, an element with atomic weighter between azinc and arsenic, but with chemical properties of those to ___.
A
germanium and silicon
22
Q
How is periodic table organised?
A
by the repeating pattern of reacivitiies and chemical properties
23
Q
Groups do or do not have similar physical properties
A
do not
24
Q
most metal oxides are ionic ___ and are basic or acidic?
A
solids , basic. This is because of the oxide present can easily form a hydroxide if reacted with water for example.
25
metal solids are usually ___ and ductile, while non metal solids are usually ___, but some are hard and soft.
malleable, brittle
26
non-metals usually form ___ solutions
acidic
27
number of energy levels correspond to periods or groups?
periods
28
Describe the 4 points of alkali metals
❑ Alkali metals are soft, metallic solids.
❑ They are found only in compounds in nature,
not in their elemental forms.
❑ Their reactions with water are highly
exothermic (release energy)
❑ One valence electron
29
Describe the 3 points of alkaline earth metals
❑ Alkaline earth metals have higher
densities and melting points than alkali
metals.
❑ Reactivity tends to increase as you go
down the group.
❑ 2 valence electrons
30
Describe the 3 points of halogens
Halogens
❑ The halogens are highly reactive.
❑ Reactivity tends to decrease as you go
down the group.
❑ 7 valence electrons (one less than a full
outer shell
31
Describe the 3 traits of Noble gases
❑ They are relatively unreactive.
❑ They are found as monatomic gases.
❑ Full outer shell of electrons
32
The atomic radius def'
one-half of
the distance between
covalently bonded
nuclei.
33
In a many electron atom, each electron is ___ to the
positively charged nucleus and ___ by the other negatively
charged electrons.
❑ The presence of other electrons around a nucleus ___ an
electron from the full charge of the ___
attracted, repelled, screens, nucleus
34
To calculate the attraction that an electron experiences due to
the nucleus and the inner shell electrons is
𝑍𝑒𝑓𝑓 = 𝑍 − 𝑆
❑ 𝑍𝑒𝑓𝑓 is the effective nuclear charge
❑ Z is the atomic number
❑ S is the screening constant (the number of core electrons
35
If there's 3 protons in the nucleus, and 3 electrons in the orbitals, what's the net charge/effective nuclear charge of the third electron?
1+ (3 -2 electrons)
36
Coulomb's law states...
the attraction of an electron to a nucleus depends only on 3 factors: the charge of the nucleus (+Z), the charge of the electron, and the distance between the two (r)
37
The ability of an electron to get close to the nucleus is...
penetration
38
Atomic radius trend
…decrease from left to right across
a period
(due to increasing Zeff).
…increase from top to bottom of a
group/column
(due to increasing value energy levels)
39
_____ the amount of energy required to remove an
electron from an atom
Ionization energy
40
Trends of ionization energy
decrease down the groups/colour, increases across the period (left to right).
41
The Nuclear charge and ionization energy relationship is ___, meaning they both ___ across the period
directly proportional , increase
42
Electron affinity def' & relationship with exothermic and endothermic
the energy change accompanying the addition of an electron to gaseous atom.
If energy is released, the electron
affinity is negative (exothermic)
If energy is absorbed, the electron
affinity is positive
(endothermic)
Electron affinity
becomes more exothermic as
you go from left to right
across a row.
43
What is 1 u equal to in g?
1 𝑢 = 1.66054 × 10−24𝑔
44
Atomic masses unit is defined by assigning 6
12𝐶 isotope of ___ a mass
of ___
carbon, 12u
45
Four major types of bonds are...
❑ Ionic bonds
❑ Covalent bonds
❑ Metallic bonds
❑ Intermolecular (van der Waals) bonds
46
Van der Waals forces def'
a general term used to define the attraction of intermolecular forces between molecules
47
what are the two types of van der Waals forces
weak London Dispersion Forces and stronger dipole-dipole forces.
48
ionic bonds def'
electrostatic attraction between ions
49
covalent bonds def'
sharing of electrons
50
metallic bonds
a strong bond formed by the attraction between positively charged metal ions and a “sea” of free-moving electrons around them. Metals lose valence electrons, hence "sea" of e-. Between two metals, strong attraction between positive metal ions and negative electron clouds.
51
Why are metals malleable, conductive, lustrous, and high melting points?
They don't have a particular orientation in space because the electrons are free to move, the ions aren't held in a lattice formation. This is why they are malleable.
| **Conductive** | Free electrons move and carry charge |
52
Why are metals shiny?
The free electrons can absorb photons in the "sea," so metals are opaque-looking. Electrons on the surface can bounce back light at the same frequency that the light hits the surface, therefore the metal appears to be shiny.
53
why do metals have high boiling temp?
Heat capacity: This is explained by the ability of free electrons to move about the solid. It's also because of their strong attrative forces.
54
intermolecular forces
are the attractive or repulsive forces between molecules
55
cation def'
a positively charged ion
56
True or false ionic compounds exist as molec, and the individual units are referred to as formula units?
false.
they don't exist as molecs, but they're individual units are called formula units
57
True or false: in an ionic compound, the positive and negative charges cancel out for a neutral compound.
True
58
Do Covalent bonds have electrostatic forces?
yes, there's...
❑ Attractions between electrons and nuclei
❑ Repulsions between electrons
❑ Repulsions between nuclei
59
Do atoms in bond share electrons equally?
Usually no, this only happens between two of the same element
60
how does electronegativity affect a covalent bond?
the electrons are held closer to the atom of higher electronegativity, making one side of the bond negative and another positive.
61
___ have electrons shared equally.
non-polar covalent bonds
62
Polar covalent bonds have electrons shared ___.
unevenly/unequally
63
Who developed a scale of electronegativities that run from a low ___ for metals in Group I to __ for fuorine.
Linus Pauling
0.7, 4.0
64
What formula to find the type of bond?
❑ Electronegativity difference: EN = higher EN – lower EN
65
Classification of bonds base on
➢ Covalent EN < 0.5
❑ If two atoms sharing a pair of electrons have equal or nearly equal
values for electronegativity the bond is nonpolar.
➢ Polar covalent 0.5 < EN < 1.7
❑ As the difference in electronegativity increases the polarity of the
bond increases.
➢ Ionic 1.7 < EN
66
polarity def'
Polarity refers to the distribution of electrical charge in a molecule or bond. It tells us if one side of a molecule is more positive or negative than the other.
67
If two atoms sharing a pair of electrons have equal or nearly equal
values for electronegativity the bond is ___.
non-polar
68
❑ As the difference in electronegativity increases the ___ of the bond __.
polarity, increases
69
what are three way to represent ionic bonding?
bhor diagram (with octets), lewis structure (with valence electrons), structural diagram (with [Cl]^-)
70
The ionic compounds is electrically ___.
neutral
71
Properties of ionic compounds
1. Crystalline solid at room temperature
2. Relatively high melting and boiling point
3. Most have high solubility in water
4. Conduct electricity as molten liquids, but not as solid (ions move independently)
5. Conduct electricity when dissolved in water (ions move freely in water)
72
Metals bond because electrons are ___ through the solid
delocalised
73
Metallic bonding traits
❑ Metals are not covalently bonded
nor are they held together by the
usual intermolecular forces.
❑ Metal bond because valence
electrons are delocalized
throughout the solid.
❑ Metallic bonding does not have a
particular orientation in space.
❑ Metal ions are not held in a rigid
lattice formation.
74
___ requires electrons to be shared equally.
pure covalent or non-polar (they are not the same)
75
Daitomic elements def' and acronym
elements that bond to each other. HOFBrINCl
76
____ forces hold the atoms of molecule together
intramolecular (they're forces that act within molec.) Intermolecular forces are BETWEEN two molecules
77
When intramolecular bonds break what happens?
chemical reaction occur
78
when intermolecular bonds break, what happens?
the state phase changes.
79
covalen compouns have __ melting and boiling points
low (meaning the covalent bonds are weaker)
80
Which has higher melting and boiling points, Non-polar or polar molecs?
Polar molec, because they both attractive
and repulsive intermolecular interactions, but
generally the attractive interactions dominat
81
What's the different between pure covalent and non-polar covalent bonds?
pure covalent only occurs between same atoms whils non-polar don't. They both share electronsequally or nearly equally. Pure covalent has a electronegativity diff. of 0 while non-polar has >0.
82
which intermolecular force is weakest? Why?
London dispersion forces are weakest because they are the result of temporary and instantaneous dipoles. This happens because at any given moment the electron distribution may be uneven, resulting in an instantaneous dipole.
83
nuclear fission occurs when...
a high unstable isotope splits into smaller particles
84
Non-polar molecule
when a molecular has no positive or negative pole
85
Is it possible for a non-polar molecule to contain polar bonds?
yes, an example is carbon dioxide. The carbon ad oxygen have different electronegativities causing uneven sharing of electrons. However, the shape of the molecule causes the electron to go to opposite directions, thus the dipoles cancel out.
86
Shape of molecules is determined by the valence shell electron domains that are ___ and ___.
bonding and non-bonding
87
VSEPR theory states
electron pairs located in bonds and lone pairs repel each other and will therefore adopt the geometry that places electron pairs as far apart from each other as possible.
88
What does VSEPR stand for?
Valence Shell Electron Repulsion
89
What are the different molecular geometries with four electron domains?
tetrahedral (all bonding) , pryimidal (one non-bonding), and bent (two non-bonding and two bonding)
90
What are the molecular geometries with three electron domains?
trigonal planar (all bonding), bent ( all three bonding)
91
What are the molecular geometries with two electron domains?
linear ( two bonding)
92
It is generally possible to remove the water of hydration by heating the hydrate.
The residue obtained after heating, called the ___ compound
anhydrous
93
In a container, a solid LiH2PO4 is present, what is the name of the coumpound?
Potassium dihydrogen phosphate
94
___ compounds contain a sulfur atom substituted for a oxygen atom in a polyatomic group
thio
95
K2S2O3 is called what?
potassium thiosulfate
96
what is the name of AlSPO3?
aluminum thiophosphate
97
___ are binary chemical compounds that have an additional oxygen atom
peroxides
98
