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MCAT GC WORKING > Periodic Table > Flashcards

Periodic Table Flashcards

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MCAT flashcards
1
Q
A
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2
Q

Where are the Alkali Metals located on the Periodic Table?

A

Alkali metals make up the first column of the Periodic Table (Group IA).

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3
Q

What is the valence shell configuration of all Alkali Metals?

What oxidation state do they ionize to?

A

All Alkali Metals have an s1 valence shell configuration.

Alkali Metals are relatively electropositive, so they will lose that 1 valence electron easily to acquire a +1 oxidation state.

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4
Q

Where are the Alkaline Earth Metals located on the Periodic Table?

A

Alkaline Earth Metals make up the second column of the Periodic Table (Group IIA).

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5
Q

What is the valence shell configuration of all Alkaline Earth Metals?

What oxidation state do they ionize to?

A

All Alkaline Metals have an s2 valence shell configuration.

They are relatively electropositive, so they will lose those 2 valence electron easily to take on a +2 oxidation state.

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6
Q

Where are the Halogens located on the Periodic Table?

A

The Halogens make up the fifth column of the p block (Group VIIA).

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7
Q

What is the valence shell configuration of all Halogens?

What oxidation state do they ionize to?

A

All Halogens have an s2p5 valence shell configuration.

They are quite electronegative, so they will accept one additional valence electron to take on a -1 oxidation state.

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8
Q

Where are the Noble Gases located on the Periodic Table?

A

The Noble Gases make up the sixth column of the p block (Group VIIIA).

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9
Q

What is the valence shell configuration of all Noble Gases?

What oxidation state do they ionize to?

A

All Noble Gases have an s2p6 valence shell configuration.

Trick question! Since they already have a completely filled octet, Noble Gases do not ionize, and they typically exist in the 0 oxidation state as free particles.

Exception: Kr and Xe, being below the 3rd row, can exceed their octet and make coordinate covalently bonded compounds such as XeF6.

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10
Q

What is the Oxygen Group, and where is it located on the Periodic Table?

A

The Oxygen Group is the group (column) below Oxygen on the Periodic Table.

It includes elements such as S and Se that are chemically similar to oxygen.

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11
Q

Where are the Transition Metals located on the Periodic Table?

A

The Transition Metals make up the entire d block.

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12
Q

Why do transition metals have high conductivity?

A

Transition metals have high conductivity due to their unfilled d subshells.

d electrons, by their nature, are loosely bound to the atom. As such, elements with partially-filled d subshells can be thought of as nuclei floating in a sea of unattached electrons, prime conditions for electrical conductivity.

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13
Q

What are the representative elements, and where are they located on the Periodic Table?

A

Representative elements are the most common elements in the solar system and the universe.

They are found in the s block and the p block of the Periodic Table.

By standard nomenclature, these are groups IA, IIA, IIIA, IVA, VA, VIA, VIIA, and VIIIA.

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14
Q

What is the valence subshell for the elements in the first two columns of the periodic table?

A

The elements of the first two columns have a valence s subshell.

Group IA has an s1 valence configuration, while IIA is s2.

Note: Helium also has a valence s subshell, but is typically listed on the farthest column with the Noble Gases, as it is chemically more similar to them than the Alkaline Earth Metals.

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15
Q

What is the valence subshell for the elements in the last six columns of the periodic table?

A

The elements of the last six columns have a valence p subshell.

Ex: Group IIIA has an s2p1 valence configuration, while VIIIA is s2p6.

Note: Although Helium is typically listed on the farthest column with the Noble Gases in VIIIA, it actually has a valence s subshell.

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16
Q

Describe the properties of metals in terms of:

  • Position in the Periodic Table
  • Electronegativity
  • Preferred Oxidation State
A

Metals are generally:

  • found in the lower-left areas of the Periodic Table.
  • low in electronegativity, losing electron density when bonded to nonmetals.
  • found in positive oxidation states when in compounds.
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17
Q

Describe the properties of nonmetals in terms of:

  • Position in the Periodic Table
  • Electronegativity
  • Preferred Oxidation State
A

Nonmetals are generally:

  • found in the upper-right areas of the Periodic Table.
  • high in electronegativity, gaining electron density when bonded to metals.
  • found in negative oxidation states when in compounds.
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18
Q

What are the main physical properties of metals?

A

Metals generally are/have:

  • good conductors of heat and electricity.
  • malleable, ductile, lustrous, and dense solids at room temp.
  • fairly high melting and boiling points.
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19
Q

What are the main physical properties of nonmetals?

A

Nonmetals are/have:

  • poor conductors of heat and electricity.
  • dull and brittle if they form solids at room temperature.
  • significantly lower melting and boiling points than metals (carbon is the primary exception).
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22
Q

How many valence electrons does Oxygen, element 8, have?

A

Oxygen has 6 valence electrons.

Oxygen is the 6th element in its row. Its valence shell configuration is 2s22p4, for a total of 6 valence electrons.

22
Q

How many valence electrons does Iron, element 26, have?

A

Iron has 8 valence electrons.

Iron is the 8th element in its row. Its valence shell configuration is 4s23d6, for a total of 8 valence electrons.

23
Q

Define:

First Ionization Energy

A

The energy required to remove one valence electron from an atom in the gas phase.

The generic ionization energy equation is:

X(g) ⇒ X+(g) + e-

23
Q

Describe the general trend of Ionization Energy heading across a row of the Periodic Table.

A

Ionization Energy increases from left to right across a row of the Periodic Table.

Other notes about Ionization Energy:

  • Atoms with fully-filled subshell will have high Ionization Energies.
  • Atoms with half-filled subshells will have higher Ionization Energies than their neighbors.
  • The Alkali and Alkaline Earth Metals have very low Ionization Energies.
24
Q

Describe the general trend of ionization energy heading down a column of the Periodic Table.

A

Ionization energy decreases heading down a column of the Periodic Table.

The further down a column an element lies, the higher the value of n for its valence electrons. Higher n electrons sit further from the atomic nucleus, and are therefore less bound to the nucleus and easier to remove.

25
Q

Which has a higher first Ionization Energy, Cl or Br?

A

Cl

Remember: Ionization Energy decreases going down a column, and Br is below Cl in the Halogens column.

26
Q

Which has a higher first Ionization Energy, Si or P?

A

P

Remember: Ionization Energy increases from left-to-right across a column, and P is to the right of Si in the Period 3.

27
Q

Define:

What is an atom’s Second Ionization Energy?

A

The energy required to remove a second valence electron from a singly-charged ion in the gas phase.

The generic Second Ionization Energy equation is:

X+(g) ⇒ X2+(g) + e-

28
Q

What are the relative magnitudes of any atom’s First and Second Ionization Energies?

A

The Second Ionization Energy is always larger in magnitude than the First Ionization Energy for every atom.

The removal of the first electron reduces the electron-electron repulsion energy of the molecule, allowing the positive nucleus to attract the remaining electrons more strongly, increasing the energy needed to remove subsequent electrons.

29
Q

How does atomic radius vary as atomic number increases across a row of the Periodic Table?

A

Atomic radius decreases across a row from left to right on the Periodic Table.

This is due to Effective Nuclear Charge. As Zeff increases, the nucleus binds the electrons more tightly, pulling them in closer.

Note: the highlighted elements represent one full row of the Periodic Table.

30
Q

How does atomic radius vary as atomic shell increases down a column of the Periodic Table?

A

Atomic radius increases down a column of the Periodic Table.

Each increasing shell can be thought of as another “layer” of electrons, outside the previous layer, increasing the atom’s size.

Note: The highlighted elements represent one column of the Periodic Table.

31
Q

Which has a larger radius, Cl or Br?

A

Br

Remember: Atomic radius increases going down a column, and Br is below Cl in the Halogens column. The valence electrons from Br are in the n=4 shell, those from Cl are in the n=3 shell.

32
Q

Which has a larger radius, Si or P?

A

Si

Remember: Atomic radius decreases from left-to-right across a column, and P is to the right of Si in the Period 3.

33
Q

Define:

Electron Affinity

A

The energy released when one valence electron is added to an atom in the gas phase.

The generic electron affinity equation is:

X(g) + e- ⇒ X-(g)

34
Q

Describe the general trend of Electron Affinity heading across a row of the Periodic Table.

A

Electron Affinity increases from left to right across a row of the Periodic Table.

The smaller an atom is, the closer a newly-added valence electron gets to the postiively-charged nucleus, and the more energy released when that electron is added.

35
Q

Describe the general trend of Electron Affinity heading down a column of the Periodic Table.

A

Electron Affinity decreases heading down a column of the Periodic Table.

The further down a column an element lies, the higher the value of n for its valence electrons. Higher n electrons sit further from the atomic nucleus, and so are less bound to the nucleus and release less energy when added.

36
Q

Define:

Electronegativity

A

An atomic property describing that atom’s tendency to attract electron density towards itself through a chemical bond.

Note: electronegativity only applies to atoms in a bond. There is no such concept as the electronegativity of a bare atomic species.

37
Q

Describe the general trend of Electronegativity heading across a row of the Periodic Table.

A

Electronegativity increases from left to right across a row of the Periodic Table.

The smaller an atom is, the closer the electrons in its bonds get to the postiively-charged nucleus, and the more strongly the electrons are attracted to the nucleus.

38
Q

Describe the general trend of Electronegativity heading down a column of the Periodic Table.

A

Electronegativity decreases heading down a column of the Periodic Table.

The further down a column an element lies, the larger its radius. This puts more space between the positively-charged nucleus and the electrons in any bonds it makes, reducing the attraction between the nucleus and the electrons.

39
Q

Which element has the highest Electronegativity?

A

Fluorine

Electronegativity increases the further right and the closer to the top of the Periodic Table an element is. Fluorine, which is the top-right element that isn’t a Noble Gas, therefore has the highest value of Electronegativity.

On the Pauling Scale, the most commonly-used scale for determining Electronegativity values, Fluorine has a value of 4.0.

40
Q

What class of elements have the highest values of Electronegativity?

A

Nonmetals are the most electronegative elements.

Fluorine is the element with the highest electronegativity, and as a general rule, the closer an element is to Fluorine, the higher its electronegativity.

Some other notes about electronegativity:

  • The Halogens are the most electronegative group.
  • Noble Gases capable of making bonds (Xe and Kr) are quite electronegative.
  • Metals, particularly the Alkali and Alkali Earth Metals, are generally quite electropositive.
41
Q

Which has a higher Electronegativity, Cl or Br?

A

Cl

Remember: Electronegativity decreases going down a column, and Br is below Cl in the Halogens column.

42
Q

Which has a higher Electronegativity, Si or P?

A

P

Remember: Electronegativity increases from left-to-right across a column, and P is to the right of Si in the Period 3.

43
Q

Describe the bond formed between two nonmetallic elements?

A

Nonmetallic elements form covalent bonds.

Nonmetals tend to have similar values of Electronegativity, and thus share electron density fairly evenly when bound together.

Covalent bonds can be purely covalent, if made between equivalent atoms (such as in O2 or in Br2). If made between non-equivalent atoms, the bond is referred to as polar covalent. In this case, the more electronegative element receives more of the electron density and carries a partial negative charge, while the less electronegative atom carries a partial positive charge.

44
Q

Describe the bond formed between a metallic element and a nonmetallic element.

A

Metallic and nonmetallic elements form ionic bonds when bound together.

In an ionic bond, the more electronegative nonmetal completely withdraws the electrons from the less electronegative metal, leaving the nonmetal a full anion (negatively charged), and the metal a full cation (positively charged).

45
Q

Describe the bonding for metals in their standard states.

A

A lattice of positively charged nuclei in a background “sea” of free-flowing negatively charged electrons.

The valence electrons of metals are loosely bound to the atomic cores, and can be considered to be effectively unattached.

46
Q

Describe the bonds formed between nonmetals.

A

Nonmetals form covalent bonds.

These can be purely covalent, such as in O2, or polar covalent, such as in H2O.

47
Q

Describe the bond formed between a metal and a nonmetal in a compound.

A

Metals and nonmetals form ionic bonds when in compounds together.

The electronegative nonmetal withdraws one or more electrons from the electropositive metal, leaving each species ionically charged.

48
Q

What are Valence Electrons, and how do they affect an atom’s chemistry?

A

Valence Electrons are the electrons in the atom’s highest energy subshells. They are chemically relevant because they are the electrons that form chemical bonds.

To find the number of valence electrons, simply count across from the left edge of the Periodic Table to the element in question.

MCAT GC WORKING (11 decks)

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