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Chemistry 135 > Unit 2 > Flashcards

Unit 2 Flashcards

(48 cards)

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1
Q

Why do bonds form?

A

The molecule has a lower energy - becoming more stable then separated atoms.
Ionic - ions formed
Covalent - electrons shared

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2
Q

Ionic Bonding

A

Electrostatic forces of attraction between oppositely charged ions
Ionic solids form because oppositely charged ions are attracted to each other in all directions. - non-directional bond
Difference in Electronegativity

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3
Q

Electronegativity

A

The ability of an atom in a molecule to attract electrons toward itself, Increasing effective nuclear charge, you can more easily attract electrons to the nucleus
Most electronegative element
- Down a column, atoms get less electronegative. Atoms are getting bigger, adding more orbitals. The attraction will be little.

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4
Q

Lewis Structure Bonding

A

Total valence electrons, draw covalent bond (2 electrons) between atoms and central atom.. Give out remaining electronegative atoms as lone pairs, if not full form multiple bonds.

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5
Q

Valence Shell Electron Pair Repulsion Model

A

Shapes of molecules arise from electron groups arranging them as far apart from others to minimize repulsion

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6
Q

Linear

A

Not counting these lone pairs because we are looking at the CENTRAL ATOM
2 electron groups around central atoms

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7
Q

Trigonal planar

A

More than one shape.
Around carbon - trigonal planar
Around nitrogen - trigonal planar

The double bond pushes the angles (not exactly 120 degrees)
TRIGONAL PLANAR - 3 electron groups around central atom

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8
Q

Tetrahedral

A

Trigonal pyramidal or tetrahedral
- Either way, it has four electron groups, that lone pair is occupying one of the sights

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9
Q

Trigonal bipyramid

A

expanded octet
SF4
seesaw

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10
Q

Octahedral

A

Lone pairs still effect the shape
6 electron groups

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11
Q

Dipole Moment

A

DIPOLE MOMENTS : a measure of the separation of charge in a molecule arising from the unequal sharing of electrons in polar bonds
Symmetry - symmetrical = cancel

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12
Q

Intramolecular Forces

A

FORCES WITHIN MOLECULE

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13
Q

Intermolecular forces

A

FORCES OUTSIDE OF MOLECULES

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14
Q

Ion-dipole attraction

A

Interaction between fully charged ion and partial charges of a polar molecule, energy of attraction increases with the charge of the ion and decreases with the square of the distance between the ion and dipole
Attraction energy = increases with the charge on the ion and decreases with the square of the distance between the ion and dipole.
Need to have an ionic compound and a covalently bound compound with a dipole moment = this interaction is only in mixtures.

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15
Q

Dipole Dipole interactions

A

In order to have dipole-dipole interactions = has to be polar molecules. Attraction between opposite charges.
there are repulsions that help orient the like charges.

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16
Q

Hydrogen Bonding

A

Hydrogen bond - intermolecular force - between molecules,
Special case of dipole-dipole
Ice floating on liquid water - less dense. Ice forms crystals that are less dense then water, partially because of the strong hydrogen bond.
Electronegativity difference is great in hydrogen bonding - strongest intermolecular force

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17
Q

London Dispersion Forces

A

Molecules that have no dipole moment - ex. Neon has no dipole moment —- but, electrons can randomly move about and concentrate in one region more than another. Making a partially negative and positive charge - INSTANTANEOUS DIPOLE
Then they can concentrate in a different region. Random movement of electrons, causing dipole moment

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18
Q

Polarizable

A

Can be distorted

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19
Q

Pressure

A

Force exerted over an area, Measuring pressure - manometer (closed tube). Tube stuck into mercury filled disk - have atmospheric pressure pushing on the dish - it will rise until the force of gravity makes the pressure balanced - then gives us the height. Forces proportional to the gravity

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20
Q

Ideal gas

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A
  1. Gases made of tiny particles moving completely randomly
  2. Total volume of particles very small compared to size of container
  3. Particles do not interact with each other
  4. Particle Collisions are elastic
  5. Kinetic energy increases with temperature
21
Q

Pressure resulting

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A

Results from gas particles colliding with container walls

22
Q

Charles Law

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A

At same P and fixed n, Volume is proportional to the temperature
V=k2T
Increasing temp - molecules move faster. Gas exerts more pressure on the lid. The lid moves out until gas is balanced with the external pressure. Keeping pressure constant - allowing volume to move

23
Q

Effusion

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A

Escape of a gas through a hole into a vaccum

24
Q

Diffusion

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A

movement of one gas through another

25
Crystalline solids
Well ordered matter within the solid - arrangement of atoms in the solid repeats itself. Very well ordered at the molecular level
26
Molecular Solids
Held together by intermolecular forces - low melt points
27
Covalent network solids
Extended structures of atoms held together by covalent bonds - very high melting points
28
Allotropes
- different structural forms of an element One super molecule - lots of covalent bonds Have to chop up covalent bonds to melt…
29
Metallic Solids
Metallic bonding between atoms - metal atoms as cations in sea of delocalized electrons, high electrical conductivity, malleable
30
Ionic Solids
Held together by electrostatic attraction between cations and anions - rock salt structure
31
Liquids
Most liquid are molecular - IMF keep particles close, but they are not strong enough, particles move past each other. Have some special properties for liquids Surface tension - amount of energy required to expand a liquid surface.
32
IMF stabilizing in liquids
water molecule in the middle has more stabilizing interactions. Lower the energy when the surface area is smaller. The stronger the forces between particles in a liquid, the greater the surface tension.
33
Capillary action
rising of a liquid in a narrow space against the pull of gravity
34
Cohesive forces
Forces between molecules
35
Adhesive forces
Forces between molecules and container walls
36
states
Solid to liquid - melting Liquid to gas - vaporization Gas to liquid - condensation Liquid to sold - freezing Gas to solid -- deposition Solid to gas - sublimation
37
Breaking attractive forces
endothermic ADDING HEAT - INCREASE IN KE - OVERCOME ATTRACTIVE FORCES
38
Forming attractive forces
releases energy, exothermic
39
Water changing state
Start at -40 degrees. Adding heat, increasing kinetic energy. At 0 degrees we reach a plateau - mix of solid and liquid, energy goes to overcoming IMF NOT BREAKING BONDS - water molecule isn't changing (heat of fusion) Stays at plateau until completely liquid. Adding heat, liquid water KE increases, molecules moving faster. Plateau at 100, heat of vaporization, then it goes into a gas (all water evaporated)
40
Thermodynamics
want to know when a phase change is a spontaneous process To know if something is spontaneous, we need to know about the free energy = negative for spontaneous
41
Gas - liquid equilibrium
Air in vacuum, below boiling point some molecules will vaporize. Vaporization initially, but when the molecules are In the gas phase, condensation can occur. Condense and vaporize at the same state - equilibrium. Leveled off pressure At 25 degrees - below boiling point, some is evaporating Vapour pressure depends on temperature - increase in temp, liquid has more KE, more molecules can escape. Gas sample has a distribution of KE Summing at a given temperature That KE distribution depends on temperature Liquid - molecules are moving, they need enough KE to move to different state (escaping the liquid)
42
Solutions
One or more substances (solutes) mixed at the molecular level (dissolved) in a medium (solvent (usually liquid)). Many reactions take place in solution
43
Forming a solution
For substances to form a solution, the solute-solvent interactions has to overcome both solute-solute interactions and solvent-solvent interactions Substances with similar intermolecular forces form solutions - like dissolves like
44
Freezing point depression
solutes lower freezing point of the solvent
45
boiling point elevation
solutes raise boiling point of the solvent
46
Vapour pressure lowering
solutes decrease vapour pressure of the solvent Non volatile solutes decrease vapour pressure of solvent (won't contribute to vaporization) - raoults law - mole fraction of solvent times vapour pressure of pure solvent equals the vapour pressure of solvent Volatile Solutes - both components of mixture contribute to vapour pressure, have to find Pressure of each then add them
47
Osmotic pressure
pressure that must be applied to a solution to prevent osmosis from a sample of pure solvent
48
Colligative properties
solution properties depend on concentration of solute, not its identity - need total concentration of everything in solution II=MRT

Chemistry 135 (8 decks)

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