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Plants > Chem 3 > Flashcards

Chem 3 Flashcards

1
Q

Gram atomic weight

A

Number of grams with 6.022 x 10^23 atoms

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2
Q

Atomic weights

A

All relative to a single 12/6C atom

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3
Q

Valence electrons

A

Electrons in outermost shell

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4
Q

Atomic number

A

Number of protons in an element

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5
Q

Alkali metals

A

Group 1A (except H), reacts vigorously with water to form strong alkaline solutions

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6
Q

Alkaline earth

A

Group IIA, oxides of these metals (chemical compounds of the metals and oxygen) form alkaline solutions in water

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7
Q

Halogens

A

‘Salt formers’, combine with metals to forms salts

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8
Q

Mass number

A

Protons+neutrons

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9
Q

Shells

Sub shells

A

‘Apartment floors’ for electrons, 1 - 7

Each floor has one or more apartments of four sizes, s, p, d, and f

Each room in a sub shell is called an orbital.

S sub shell - single room, one orbital that will hold a max of 2 electrons

P sub shell - three rooms, 6 electrons

D sub shell - five rooms, 10 electrons

F sub shell - seven rooms, 14 electrons

  • electrons prefer lowest floor, and prefer to live one to a room until each room in apartment has one occupant. They then will put up. Each apartment can only hold two

** shells 4-7 each have four sub shells

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10
Q

Principle of maximum multiplicity

A

Electrons occupy as many orbitals as possible in the same sub shell before pairing with another electron

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11
Q

What explains the similar chemical properties of elements within the same group?

A

Each group of elements in the periodic table has similar sub shells with similar numbers of electron

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12
Q

Why does an atom gain or lose electrons

A

General rule - atoms tend to form sets of eight electrons in their outer most shell (valence electrons) when they combine to form compounds. Atoms with six or seven valence electrons have a great attraction for one or two extra electrons, and atoms with only one, two or three valence electrons have only a weak hold on those electrons. An atom with six or seven valence electrons will tend to gain electrons from another atom with one, two or three valence electrons when those atoms combine chemically

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13
Q

Octet rule

A

The general tendency of atoms to form sets of eight electrons in their outer most shells when they combine to form compounds.

Generally tend to form electron configurations that are similar to noble gases.

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14
Q

Electrostatic bond

A

Force that holds Negative and positive ions together in an ionic bond.

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15
Q

Formula for ionic compound

A

Most metallic element is written first

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16
Q

Covalent bond

A

One or more valence electrons from each atom are shared. The resultant compound is made up of molecules, not ions. I

Normally formed between non metals.

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17
Q

Ionic bond

A

Elements held together by electrostatic force.

Normally formed between a metal and nonmetal.

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18
Q

Coordinate covalent bond

A

Both electrons in a shared pair of electrons come from the same electron, otherwise same as normal covalent bond

Examples include ammonia NH3, boron triflouride BF3

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19
Q

Electronegativity

A

The attraction an atom has for the electrons in a covalent bond. Electrons will spend more time around the element with the stronger bond. This will create a slight + and - bond on different sides of the compound, called a polar bond.

Can classify bond types for compounds with two atoms based on electronegativity difference-

Covalent, less than .5
Polar Covalent, between .5 and 1.5
Ionic, greater than 1.5

Only time 100% covalent is when it is equaled to zero, which happens with two of the same atoms

Liquified ionic compound will conduct electricity, but covalent will not. A polar compound will be attracted to charge rod, a no polar will not

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20
Q

Compound shapes (covered by book)

A

Linear, straight line can pass through nuclei. Often non polar if atoms of same element cancel each other out (example of CO2, where Os are side by side with C atom)

Bent,

  • planar, atoms fall within the same flat surface (atoms on this page)
  • nonplanar, one or more atoms fall outside for surface (above or below page)
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21
Q

Ionization energy

A

Energy required to remove electrons from a neutral atom, and increases as you move up and right in the periodic table. You would say He requires more ionization energy than Na.

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22
Q

Molecule of

Element

Compound

A
  • consists of one or more atoms of the same kind. A single atom of an element can be referred to as an atom or molecule
  • usually involves two or more atoms combined with a covalent bond

-

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23
Q

Molecular weight

Formula weight

A

Molecular
Sum (Element amu x number of elements)

Formula
Same equation as molecular, but use for either molecular or ionic compounds

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24
Q

Chemical formula

Empirical formula

A

Molecular
Describes the actual number of atoms contained in each molecule. Must know relative weights of the elements in a compound.

Empirical
Simplest possible chemical formula that represents the ratios of the atoms within an unknown molecule

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25
Q

Mole (mol)

A

6.022 x 10^23 units

Can say mole of atoms, electrons and ions

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26
Q

Binary compounds

A

Containing only two electrons, typically a metal and non metal

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27
Q

Binary compounds

A

Containing only two electrons, typically a metal and non metal

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28
Q

Oxyacids

A

Hydrogen + nonmetal + oxygen

Named according to the most commonly occurring acid, the ‘ic’ acid. ‘Ic’ is added to the name of the nonmetal

As the alphabet reaches the letter O, the number of oxygen atoms in the “ic acid increases from three to four

Acids with one less oxygen than ‘ic’ are “ous

Two less oxygen atoms = “hypo-ous

One more oxygen = “per-ic

Written as H + polyatomic acid, with necessary subscripts to make the charges zero

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29
Q

Oxyacids salts

A

Metal + nonmetal + Oxygen

Derived from oxyacids

Named by writing the name of the metal followed by the name of the polyatomic acid ion.

“Ic acids form “ate salts

“Ous acids form “ite salts

“Hypo-ous acids form “hypo-ite salts

“Per-ic acids form per-are salts

Formulas are written with the metal symbol first followed by the polyatomic acid ion, with necessary subscripts to make the charges equal zero

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30
Q

Oxidation numbers for metals with more than one

A

Iron 2 and 3+

Copper 1 and 2

Gold 1 and 3

Tin 2 and 4

Chromium 2, 3 and 6

Mercury 1 and 2

Lead 2 and 4

31
Q

Kinetic molecular theory (gas)

A
  • gases are made up of individual molecules that are in constant rapid motion (possessing kinetic motion), bouncing into one another and ceiling / walls (top and bottom equally)
  • molecules are further apart than in liquids or solids
  • temperature is measure of average kinetic motion
  • pressure represents the collision of molecules against walls of container. Measured as the number of collisions on a given area (sq unit) of the container
  • molecules occupy no volume and have no attraction to one another. Space in between molecules is what makes up volume
  • Gases can be more easily compressed than liquids and solids
32
Q

Diffusibility

- gas

A

Substances mix completely

  • gas - more rapid than in other states
33
Q

Boyles law

A

At a constant temperature, the volume of a confined dry gas is inversely proportional to the pressure

PV = a constant at a constant temperature

P1V1 = P2V2

34
Q

Torr - Atmosphere - Kilopascal

A

A pressure that will suppport a volume of Mercury 1 millimeter high.

1 atm (atmosphere) is equal to 760 torr

1 atm = 760 torr = 101.325 kilopascals (kPa)

35
Q

Absolute temperature scale

Charles law

A
  • Measure of the average kinetic energy of gases. Doubling absolute (Kelvin scale) temperature doubles the average kinetic energy and thus pressure if volume is constant, or a doubling of the volume of the pressure is constant. Kelvin zero point is -273 degrees Celsius. At this point theoretical volume of gas is zero
  • with constant pressure, volume of a gas is directly proportional to the temperature

V / T = a constant at constant pressure
Or
V1/T1 = V2/T2

36
Q

Partial pressure

Dalton’s Law of Partial Pressure

A

Each gas molecule contributes equal amounts to the total pressure regardless of what kind of molecule it is. Partial pressures are Proportional to the number of moles of the gases in a mixture

Total pressure of a mixture of gas is the sun of partial pressures

Dalton - total pressure of a mixture of gases is the sun of all of the partial pressures of the individual gases. The molecules of each gas exert the same pressure within the mixture as they would if they were not in the mixture

p gas = n gas / n T x P T

Lower case p = partial pressure
n gas = moles of the gas in question
n T (T is subscript) = total moles of all gases
Upper case P = total pressure
37
Q

Diffusion vs effusion

*Standard temperature and pressure

A

Diffusion - mix spontaneously throughout each other

Effusion - process of a gas going through a small opening into a vacuum (absence of a solid, liquid or gas)

Graham’s law can be applied to both. At a constant temperature, the rate (velocity) of effusion of a gas is inversely proportional to the square root of its formula weight. Assumes temp and pressure are same for both gases. Can be modified to deal with the density of gases.

The heavier the formula weight, the slower the rate of effusion

v1/v2= sqrt(M2/M1)

v1/v2=sqrt(d2/d1)

*to have an accurate measure of the density of a gas, we must specify a temperature and pressure. Standard temperature and pressure (STP) have been set arbitrarily at 0 degrees Celsius and 1 atm. The volume of 22.4 liters per mile at STP is a reasonable approx for real gases

38
Q

Combined gas law

Ideal gas law

A

Combined Charles and Boyle

P1V1/T1 = P2V2/T2

Simper version ‘ideal’…

PV=nRT

n=represents the number of moles of the gas
R=universal gas constant, a numerical value made up of the combination of standard pressure (1 atm or torr), standard temperature (273 K), and the fact that 1 mol of gas occupies a volume of 22.4 liters at STP.

39
Q

Atoms in a solid

A

Vibrate in place, fixed location.

Force of attraction is strong. Changes in temp and pressure only cause little change

Melting - changing from solid to liquid state

40
Q

Sublimation

A

Changing directly from solid to gas

41
Q

Crystalline solid, or crystal

A

A solid that exists in a definite three dimensional geometrical shape. Seven different basic shapes.

Crystal lattice - The arrangement of atoms within the crystal is known as the crystal lattice

Polymorphic - substance can exists in more than one crystalline form. The different crystalline forms are known as allotropes

42
Q

Amorphous

A

Solid that doesn’t have crystalline shape.

Glasses - Those solids that appear to be crystals but are not. Contrasting qualities include:

                    Crystal.                Amorphous

Melting point. Sharp. Large ranges

When cleaved. retain structure. Fracture into uneven surfaces

Bond strength. Consistent. Varies through solid

43
Q

Heat of fusion

A

Energy required to change solid to a liquid while remaining at melting point.

Calories per gram, or calories per mole

44
Q

Specific heat of a substance

A

Calories required to raise temp of 1 gram of substance by 1 degree C. Related to atomic weight

Atomic weight = 6.2 / specific heat

45
Q

Types of solids

A

Ionic - ionic bonds, hard and high melting point. Do not conduct electricity as a solid, but when liquid (molten) they do so well. Most dissolve well in water. Examples - NaCl and KCl

Covalent - covalent bonds, extending from atom to atom throughout the crystal lattice. Very strong, very hard and very high melting points. Non conductors of electricity either as solid or liquid, and not soluble in water. Examples - diamonds, silicon carbide SiC, and aluminum nitride AlN.

Molecular - made up of polar or non polar molecules. Held together by weak van der waals force. Soft, low melting points and do not conduct electricity. Examples of non polar - dry ice CO2 and naphthalene (mothballs). Examples of polar - H2O, solid ammonia NH3 and solid sulfuric acid H2SO4.

Metallic solid - made up positive ions held together by highly mobile electrons or as positive ions in a sea of electrons. Electrons move freely, making them great conductors of electricity. Hardness and melting point varies widely. Has malleability, ductilibility and luster. Examples include iron, copper and silver.

46
Q

Atomic and ionic radii

A

Typically measured in angstrom (1 A = 1 x 10^-8)

Positive ions are smaller than their parent atoms because they’ve lost one or more electrons. Protons now outnumber electrons and thus draw the remaining in closer to center. Also, loss of electrons in outer most shell reduces size.

Negative ions - Opposite process

47
Q

Isoelectric

A

Ions and atoms that have the same number of electrons. Na+ has the same number of electrons as Ne.

Includes (these have ten electrons)
N3-, O2-, F-, Ne, Na+, Mg2+, Al3+, Si4+

48
Q

Intermolecular attractive forces

A

Cause molecules to stick together and coalesce into a liquid. Strength impacts the flow of the liquid.

Viscosity measures resistance to flow. Higher resistance –> higher viscosity

49
Q

Surface tension

A

Caused by a difference in direction of intermolecular attractive forces between those molecules at the surface of a liquid and those in the body of the liquid. Strength of force depends on intermolecular force.

Those at surface are pulled down, those below have equal forces pulling all around.

Downward pull results in the liquid forming as small a surface as possible.

A droplet would be circular is not for gravity.

50
Q

Vapor pressure

A

Pressure exerted by the vapor (gaseous) phase of a liquid in a closed container.

‘Average kinetic energy’ in a liquid suggests a variation in energy. Those with higher energy will escape if they are located near the surface. Those of lower energy remain, resulting in drop of temp for those remaining.

To make up for the loss of kinetic energy, a liquid undergoing evaporation will draw heat from its surroundings.

51
Q

Vapor pressure

A

When rates of evaporation and condensation are equal. Said to be in equilibrium. These will vary based on temperature and pressure.

52
Q

Le Chatelier’s Principle

A

Governs equilibrium systems… a ‘stress’ on a system in equilibrium will cause a shift to a new equilibrium so as to relieve the stress

53
Q

Boiling point

A

Temp at which the vapor pressure of a liquid equals the surrounding pressure. The higher the temp, the greater the vapor pressure.

General rule - compounds of the same family on the periodic table increase in boiling point with increasing molecular weight, all other factors being equal.

Other factors include:
Strength of intermolecular attractiveness
-polar molecules’ attractive force (hydrogen bonding)
-van der waals
-London force

54
Q

London forces

A

Caused by electrons from one molecule affecting f the orbit of electrons in another molecule at very specific and constantly changing moments in time.

55
Q

Volatility

A

A measure of the tendency of liquids to evaporate.

56
Q

Miscibility

A

Ability of liquids to mix. Polar and non polar liquids will not mix with each other. Think oil and water.

57
Q

Heat of vaporization

A

Heat required to vaporize water at the boiling point.

Varies based on liquid.

Think refrigerant system. Take advantage of the fact that a liquid which is vaporized at its boiling point requires heat.

As substance is condensed / compressed to point of condensation, it releases heat (back of fridge). As it is vaporized, it absorbs heat from the surround area (inside fridge)

58
Q

Critical temperature

A

Temperature above which the intermolecular attractive forces can no longer overcome the kinetic energy in order to liquify a gas. At any temp above this point, a gas will not liquify no matter how much pressure is applied.

59
Q

Solute vs solvent

A
  • the one being dissolved

- the one doing the dissolving. Basically there’s more of it

60
Q

Electrolyte

A

A soluble ionic solid in aqueous solution. It conducts electricity. Ions in an ionic compound disassociate and are free to move around

Non electrolytes may dissolve but molecules stay together.

61
Q

Solute reacting in water

A

Oxides if alkali metals (Na2O, K2O, Li2O) and alkaline earth metals (MgO, CaO, BaO) react with water to form metal hydroxide solutions.

Na2O + H2O –> 2NaOH

Sodium hydroxide NaOH ionized to yield Na+ and oh-.

A compound that ionizes to form OH- in aqueous solution is called a base

62
Q

General equations for reaction of an

  • Alkali metals with water
  • alkaline earth with water
A
  • M2O + H2O – > 2MOH

- MO + H2O – > M(OH)2

63
Q

Acid

A

The oxides of a number of non metals react with water to form acids

N2O5 + H2O –> 2HNO3

Where HNO3 (H+ and NO3-) is the new solute

64
Q

Acid anhydrides

A

Nonmetal oxides, named this way because they form acids when mixed with water

65
Q

Basic anhydrides

A

Metal oxides, named this way because they form bases when added to water

66
Q

Hydrate

Water of hydrate

A

Compounds with water within the structure of their crystal lattices.

Water that is contained
-when compound is heated enough, it’ll release of the water molecules.

The dots (•) indicate that water is an integral part of the crystal lattice of the hydrate

67
Q

Hygroscopic

Deliquescent

A
  • substances that can absorb moisture from the atmosphere

- some can absorb enough to dissolve in their own water of hydration

68
Q

Solution that is

Saturated

Unsaturated

Supersaturated

Concentrated

Dilute

A
  • solution containing max amount of solvent at a given temperature
  • contains less than maximum at a give temperature
  • solution holds more solute than is normal, done by adding at higher temp and then cooling
  • solution with relatively high amount of solvent
  • solution with relatively low amount of solvent
69
Q

Molarity

Vs Molality

A
  • Indicates the moles of solute per liter of solution. Specifies concentration. Use when dealing with the concentration of the solute in the solutions
  • indicates the moles of the solute in each kilogram of solvent. Best when looking at colligative properties. (Volumes would have to be converted to weight… need density to calculate that).
70
Q

Colligative properties

A

Properties (freezing / boiling points, vapor pressure) that vary according to the ratio of the weights of the solute and solvent

71
Q

Mole fraction and vapor pressure

A

Placing a liquid solute in a liquid solvent lowers the vapor pressure of the solvent. Amount contributed to pressure by either solvent or solute depends on the mole fraction (number of moles of one of the components of the solution (either solute or solvent) divided by the total number of moles in the solution.this leads to Raoult’s Law (frame 75).

Vapor Pressure contributed by liquid A in the solution = mole fraction of liquid A in the solution x vapor pressure of liquid A in its pure form

72
Q

Molal boiling point elevation

A

The presence of 1 mol of any nonvolatile non electrolyte such as sugar added to 1 kg of water raises the boiling point of the water by .51 degrees celcius.

In other words, if a nonvolatile solute is added to a solvent, the boiling poi t of the solvent is raised because the vapor pressure is reduced.

Predicting boiling point:

Increase in boiling point = the boiling point constant x molality of the solutions .

73
Q

Freezing point depression

A

A nonvolatile solute lowers the freezing point of the solution.

For each mole of a nonvolatile non electrolyte added, in 1 kg of water, will lower freezing point by 1.86 degrees celcius.

*each mole has a specific number of particles, so we can say that the greater the number of solute particles in a solvent the greater will be the freezing point decrease and boiling point increase

74
Q

Endothermic

Exothermic

A

Chemical reaction requiring heat to proceed
-increasing temp will increase the concentration of the products

Chemical reaction giving off heat during a reaction
-opposite

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