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Chemistry > Unit 4: The Periodic Table > Flashcards

Unit 4: The Periodic Table Flashcards

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1
Q

electronegativity

A

tendency for the atoms of the elements to attract electrons during chem. reactions.

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2
Q

the pauling scale relies on what two factors?

A

ionization energy and electron affinity

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3
Q

ionization energy

A

the amount of energy needed to remove the highest energy electron from an element in its GASEOUS PHASE

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4
Q

electron affinity

A

the change in energy when an electron is added.

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5
Q

electroegativity trends in groups

A

electronegativity decreases down a group, because PEL’s are added going down a group which creates a greater atomic radius, decreasing the nuclear attraction that would increase electronegativity.

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6
Q

shielding

A

the buffering of attraction between the positive nucleus and the outermost electron shell.

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7
Q

electronegativity with noble gases

A

they do not have electronegativity values because they were historically thought to not bond, and even though they sometimes do, no values were assigned anyways.

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8
Q

electronegativity trends down a period

A

from left to right, electronegativity increases, since the valence electrons become more full, making it easier to gain electrons to fill up the shell, and atomic radius decreases as the nuclear attraction increases.

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9
Q

What is the element with the lowest electronegativity?

A

cesium, with only 1 valence electron and the largest atomic radius

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10
Q

what is the element with the highest electronegativity

A

fluorine, small atomic radius and high # of valence electrons.

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11
Q

cation

A

atom that loses electrons, a positively charged ion.

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12
Q

what kind of elements form cations?

A

metals since they have leses valence electrons, easier to lose than gain to create a full octet/duet.

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13
Q

Is a cations radius smaller or larger than the original atoms?

A

the radius is smaller, since the loss of the electron results in less shielding, making the attraction from the nucleus pull the electrons closer to the atom.

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14
Q

anion

A

atom that has gained an electron, negatively charged

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15
Q

what kinds of elements form anions?

A

nonmetallic anions, bc they have high ionization energies and readily gain electrons

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16
Q

is the radius of an anion larger or smaller than that of the original atoms?

A

larger, because the addition of the electron makes the atom bulge due to electron-electron repulsion

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17
Q

metalliods/semimetals

A

elements with properties between or in combination of those of nonmetals and metals.

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18
Q

what are the six metalloids?

A

boron, silicon, germanium, arsenic, antimony, tellurium BATGAS

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19
Q

what is the only metal that is liquid at STP?

A

mercury

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20
Q

what is the only nonmetallic element that is liquid at STP?

A

bromine

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21
Q

what are the only two elements on the periodic table that are liquid at STP?

A

bromine and mercury

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22
Q

what is considered the most metallic element, chemically?

A

cesium (since francium is radioactive) because it has only one v-electron, has the largest radius due to many PELs, which causes shielding, and the lowest ionization energy.

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23
Q

what group is the most reactive metallic elements?

A

alkali metals

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24
Q

what is considered the most chemically nonmetallic element?

A

fluorine, since it has the smallest radius of all group 17 elements and only has 1 missing v-electron, allowing it to be the most attractive element.

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25
Q

what are the most reactive non metallic elements?

A

group 17 halogens, who react with alkali metals, combining to form salts

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26
Q

which group 18 elements have some chemical reactivity?

A

Xenon and Krypton have oxidation states other than 0, though noble gases all have full valence shells and are very stable as a result, and so usually will not bond.

27
Q

name the monatomic elements

A

group 18 noble gases

28
Q

name the diatomic elements

A

H O F Br I N Cl

29
Q

name the five elements that are gases at STP besides the noble gases

A

hydrogen, nitrogen, oxygen, fluorine, and chlorine. (the diatomic elements minus iodine and boron)

30
Q

after which number element are all elements considered naturally occurring radioactive isotopes?

A

83

31
Q

allotropes

A

2 or more different molecular forms of the same element in the same physical state

32
Q

name the two allotropes of oxygen

A

ozone (O3) and oxygen gas (O2)

33
Q

name the three allotropes of carbon

A

diamond, graphite, and buckminsterfullerenes (C-60)

34
Q

atomic radius

A

half that distance between the nucleus of the 2 bonding atoms.

35
Q

measuring atomic radius

A

cannot be measured directly since atoms do not have defined boundaries, though there are many ways to estimate

36
Q

what are the two ways to estimate atomic radius?

A
  • using X-ray diffraction if the element is a solid, crystalline structure, to estimate the distance between nuclei.
  • for diatomic elements it is only possible to estimate the distance between the nuclei of the two bonded atoms.
37
Q

atomic radius trends down a group

A

atomic radius increases as you go down a group, bc more PEL’s are being added, increasing shielding as well.

38
Q

atomic radius trend across a period

A

generally decreases since the PEL remains the same while more and more v-electrons are added, increasing the NUCLEAR CHARGE while the shielding stays the same.

39
Q

in a single period, atomic radius and atomic number have an

A

INVERSE RELATIONSHIP

40
Q

when a substance is ionized it results in a…

A

1+ charge

41
Q

why are there jumps in ionization energy?

A

valence electrons are easy to take away, so they require low amounts of ionization energy to be stolen. Once they are all gone, it takes much more energy to take away a core electron, which accounts for the jump.

42
Q

what is the unit for ionization

A

kj/mol

43
Q

what element has the highest first ionization energy?

A

helium, because it has both a full electron shell and also the smallest atomic radius, which separates it from the rest of the noble gases, since a smaller radius = greater attraction = harder to steal the electron

44
Q

ionization group trends

A

as atomic # increases in group the ionization energy decreases, since the PELs increase, increasing shielding, which makes it easier to steal the electron.

45
Q

ionization trends across a period

A

first ionization energy increase across a period, since the nuclear charge increases, strengthening the hold on the electrons.

46
Q

which subatomic particle effects chemical and physical properties the most?

A

the electron

47
Q

the arrangement of the periodic table is related to

A

electron configuration

48
Q

what are the four categories electrons can be classified into based on electron configuration?

A

noble gases, rep. elements, transition metals, and inner transition metals.

49
Q

electron confg: noble gases

A

they have a full outer shell and are inert gases bc they do not bond, belonging to block p except for Helium, which is in the s block.

50
Q

electron config: representative elements

A
  • elements with outermost p and s sublevels only partially filled
  • group number tells you the number of valence electrons
  • the representative element group names are the alkali metals, alkaline earth metals, and the halogens.
51
Q

electron config: transition metals

A
  • group B elements
  • contains 3s elements
  • defined as the elements with an incomplete d orbital.
  • belongs to the d sublevel block but not all d-block elements are transition metals
  • electrons are added to the d sublevel rather than the outermost shell, disrupting the pattern
  • transition metals fill the PEL one less than the period number.
52
Q

electron config: inner transetals

A
  • the lanthanides and actinide series, belonging to the F sublevel, where electrons are added to the f shell rather than the outermost shell and the PEL added to is 2 less than the period.
53
Q

electrno config: lanthanides

A

electrons are added to 4f

54
Q

electron config: actinides

A

electrons added to 5f

55
Q

dmitri mendeleev (1843-1907)

A
  • mid 1800s arranged elements in vertical columns with increasing atomic mass, arranged by chem. and physical properties
  • blank spaces were from undiscovered elements
  • he could predict properties by comparing to the known elements.
56
Q

henry mosely (1887-1915)

A
  • determined nuclear charge of the atom, aka the atomic #

- arranged the table according to atomic # like the modern periodic table.

57
Q

periodic law

A

when elements are arranged in order of increasing atomic number, there is a periodic pattern of chemical and physical properties

58
Q

what kinds of elements form compounds and solutions with COLOR

A

transition metals

59
Q

what kinds of elements have MULTIPLE oxidation states

A

transition metals

60
Q

name the four properties of metals

A
  • malleable
  • ductile
  • luster
  • good conductors of both heat and electricity
61
Q

ductile

A

can be made into wire

62
Q

luster

A

ability to reflect light

63
Q

group 1 metals are so reactive that you will never find them _____ in nature, only in the form of _____

A

free, salts

64
Q

how do you extract the group 1 element from the compound

A

run electric current through it, Cl gas for example will evaporate and leave behind the group 1 element.

Chemistry (17 decks)

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