Download Metals

Chapter 6.1 Periodicity
Ox Quiz#1, Ox Quiz#2
now!
tomorrow!
• By now, you know that the Periodic Table is
organized by increasing atomic #, by orbitals,
and by the number of valence electrons.
• But… There’s still more!
• The organization of the periodic table can tell
you several properties of elements, if you
know what you’re looking for.
• Periodic Law: When the elements are arranged in
order of increasing atomic number, there is a
periodic pattern in their physical and chemical
properties
• Elements within the same column, or group, have
similar properties.
• Elements within the same row, or period, change
from left to right across the PT, but repeat in the next
period.
Metals
From everyday life, you should be familiar with the word “metal”
and should know of several metals… But just how many of the
elements on the periodic table are metals?
Metals
Roughly 80% of elements are metals.
• Generally, elements can be classified as being
metal, nonmetal, or metalloid.
• Each can be broken down into sub-categories,
which can be seen in the periodic table in your
book.
Metals: - Good conductors of heat/electricity.
- Ductile (can be made into a wire)
- Most are malleable (can be molded)
- Tend to be solid at room temp.
Who am I?
13p
14n
Metals: - Good conductors of heat/electricity.
- Ductile (can be made into a wire)
- Most are malleable (can be molded)
- Tend to be solid at room temp.
Al
13p
14n
13p
14n
Outer energy levels overlap
Metals: - Good conductors of heat/electricity.
- Ductile (can be made into a wire)
- Most are malleable (can be molded)
- Tend to be solid at room temp.
Nonmetals: - Mostly poor conductors of
heat/electricity.
- Tend to be gaseous at room temp.
- Not malleable
forbidden gap
16p
16n
16p
16n
Metals: - Good conductors of heat/electricity.
- Ductile (can be made into a wire)
- Most are malleable (can be molded)
- Tend to be solid at room temp.
Nonmetals: - Mostly poor conductors of
heat/electricity.
- Tend to be gaseous at room temp.
- Not malleable
Metalloids: - Have properties of both metals
and nonmetals.
- Behavior often controlled by
changing conditions
Sub-categories of metals and nonmetals:
You must memorize the names of the groups!!!
Ch6 HW#1
Chapter 6 HW#1
1. List your favorite element from the halogen family, then draw its dot diagram.
2. List your favorite element from the alkali metals, then draw its dot diagram.
3. Make a statement about the reactivity of the alkali metals and the halogens.
4. Who am I? I am the only element that falls along the metal/nonmetal
dividing line that is strictly a metal. Engineers like to make lightweight structures
out of me and my alloys. I carry 3 valence electrons…
5. Who am I? I can sometimes be classified as a metal or a metalloid,
and chemically, my valence electron configuration is similar to
oxygen, except I got a whole bunch more energy levels.
Did you know that Rutherford used one of my radioisotopes
in his famous Gold Foil Experiment.
Ch6.2 Ions and Atomic Size Ox Quiz#2 today!
By looking at the periodic table, how does the atomic mass of a halogen
compare to that of an alkali metal in the same period?
Atomic Mass increases across a period. (Ch4)
How do you think the overall size of the atom compares for atoms of the same
two elements?
Atomic Mass increases across a period. (Ch4)
Atomic Radius (size of the atom) decreases across a period.
This is because the increased number of protons (positive charges) in the
nucleus creates a greater electromagnetic pull on the electrons
(negative charges) orbiting around the outside.
Bigger
Smaller
Atomic Mass increases across a period. (Ch4)
Atomic Radius (size of the atom) decreases across a period.
This is because the increased number of protons (positive charges) in the
nucleus creates a greater electromagnetic pull on the electrons
(negative charges) orbiting around the outside.
Atomic radius does increase as you move down a group due to the fact
that there are more energy levels.
Ions
• We know that atoms have a neutral charge
because the charges of protons and electrons
cancel each other out.
• But, when atoms react with their surroundings,
they tend to either gain or lose an electron,
which leaves it with either a positive or negative
charge.
• An atom that has gone through this kind of
reaction is called an ion.
• Cation: an atom that has lost an electron and is left with an
overall positive charge.
Al:
13p
14n
• Anion: an atom that has gained an electron and is left with an
overall negative charge.
• Cation: an atom that has lost an electron and is left with an overall
positive charge.
Al:
Al3+
10 electrons
13 protons
=
13p
13p
14n
14n
• Anion: an atom that has gained an electron and is left with an overall
negative charge.
S:
16p
16n
• Cation: an atom that has lost an electron and is left with an overall
positive charge.
Al:
Al3+
10 electrons
13 protons
=
13p
13p
14n
14n
• Anion: an atom that has gained an electron and is left with an overall
negative charge.
S:
S2–
18 electrons
16 protons
16p
16p
=
16n
16n
Cation:
Anion:
• Lost an electron
• Positive charge
• Denoted by a (+) after the
symbol (Na+, Mg2+, etc.)
• Tends to occur in metals.
• Decreases the atomic radius of
the atom.
• Gained an electron
• Negative charge
• Denoted by a (–) after the
symbol (F–, S2–, etc.)
• Tends to occur in nonmetals.
• Increases the atomic radius of
the atom.
Cation:
Anion:
• Lost an electron
• Positive charge
• Denoted by a (+) after the
symbol (Na+, Mg2+, etc.)
• Tends to occur in metals.
• Decreases the atomic radius of
the atom.
• Gained an electron
• Negative charge
• Denoted by a (–) after the
symbol (F-, S2-, etc.)
• Tends to occur in nonmetals.
• Increases the atomic radius of
the atom.
Cations and Anions are both results of
an ionic reaction… For example, in
sodium chloride (table salt):
lose
cation
gain
anion
See textbook or online for ionization/atomic radius chart.
Ch6 HW#2 + Lab6.1 Pre-lab Questions
Chapter 6 HW#2
1. Explain why fluorine has a smaller atomic radius than both oxygen and chlorine.
2. Arrange these elements in order of decreasing atomic size:
sulfur, chlorine, aluminum and sodium.
Does your arrangement demonstrate a periodic or group trend?
3. Which particle has the largest radius in each atom /ion pair?
a. Na, Na+
b. S, S-2 c. I, I-1 d. Al, Al+3
4. Chlorine, selenium, and bromine are located near each other on the period table.
Which of these elements is the smallest atom?
5. Phosphorus, sulfur, and selenium are located near each other on the periodic table.
Which of these elements is the largest atom?
6. Scandium, yttrium, and lanthanum are located near each other in the periodic table.
Which of these elements is the largest atom?
7. Which of the following atoms is smallest: vanadium, chromium or tungsten?
8. Which of the following atoms is smallest: nitrogen, phosphorus or arsenic?
NaF NaCl NaBr NaI Unknown A
NaF NaCl NaBr NaI Unknown A
?
NaF NaCl NaBr NaI Unknown A
?
NaF NaCl NaBr NaI Unknown A
?
(start over)
NaF NaCl NaBr NaI Unknown A
?
No tree given for part B. Follow the instructions!
- Add H2SO4
- clean, then add NH4OH
- clean, then add …
(start over)
Mg2+ Ca2+ Sr2+ Ba2+ Unknown B
?
Ch6.3 Ionization Energy
• We know that, as electrons absorb energy, they can move up to higher energy
levels.
• If they absorb enough energy, they can leave the atom entirely.
• Ionization Energy: The energy required to remove an electron from an atom.
See text or online for a
list of first, second and
third ionization energies
for the first 20 elements.
• The first ionization energy is
the energy required to
remove one electron, leaving
a +1 charge behind.
• The second ionization energy
is the energy to remove a
second electron (once the
first has already been
removed), leaving a +2 charge
behind.
• And so on, and so on…
• In general, the ionization energy increases from left to right
across a period and from bottom to top in a group on the PT.
(Table 1 for HW)
This figure shows the trend of the first ionization
energies of the representative elements (those in the s
(Table 2 for HW)
• Graphically, we can also see the trends in ionization energies.
• Notice that the ionization energies of the alkali metals (Li, Na, K,
etc.) decrease with atomic number, as well as the noble gases (He,
Ne, Ar, etc.), just as they would in their respective rows.
(Table 3 for HW)
Ex1) Lithium has a very low 1st ionization energy. But there is a tremendous jump
to its 2nd IE. Explain.
3p
4n
Ch6 HW#3 1 – 6
Chapter 6 HW#3
1. Distinguish between the first and second ionization energies of an atom.
2. Refer to Table 1, indicate which element in each of the following pairs
has the greatest first ionization energy.
a. lithium, boron
b. magnesium, strontium
c. cesium, aluminum
3. Chlorine, selenium, and bromine are located near each other on the period table.
Which of these elements has the highest ionization energy?
4. Phosphorus, sulfur, and selenium are located near each other on the periodic table.
Which of these elements has the highest ionization energy?
5. Refer to table 2, there is a mini-spike in the graph just before atomic # 50.
What element might this be? Justify your choice.
6. Refer to table3, why does carbon have such a significant jump
in ionization energy from I4 to I5?
Ch6.4 Electronegativity
- the ability of an atom to attract electrons when the atom
is in a compound.
This property also tends to increase from left to right,
and from bottom to top.
(Noble gases are excluded from the effects of
electronegativity because they have full valence shells
and rarely form compounds)…
(EN table, refer for HW)
Atoms desire to have either empty or full valence shells
… In other words, they like to have either 0 or 8 electrons in their outer shells.
Elements with low numbers of valence electrons (alkali metals, alkaline metals, etc)
Tend to have low ionization energies/low electronegativity
(easier to achieve 0 by taking electrons away)
Elements with high numbers of valence electrons (halogens, etc)
Tend to have high ionization energies/high electrogenativity
(easier to achieve 8 by gaining electrons)
(Use this electronegativity table in for your homework.)
Covalent bonding: Occurs when there is electron sharing between
nonmetals and/or metaloids (high electronegativity).
Ionic Bonding: Occurs when there is electron transfer from a metal (low ionization
energy) to a nonmetal (high electrogenativity), forming both + and – ions.
Covalent bonding: Occurs when there is electron sharing between
nonmetals and/or metaloids (high electronegativity).
Ionic Bonding: Occurs when there is electron transfer from a metal (low ionization
energy) to a nonmetal (high electrogenativity), forming both + and – ions.
Electronegativity Difference – subtract the #’s from the EN Table.
1.67 is the magic #.
Ionic Bonds > 1.67
Covalent Bonds < 1.67
Ex1) Decide if the following elements bond covalently or ionic:
a. Na and Cl:
b. Si and O:
c. O and O:
Ch6 HW#4
Lab 6.1 Qualitative Analysis
In the laboratory, an important activity of the chemist is to determine what elements
are present in a given sample. Specific reagents are not available to identify each type of ion
or atom. Therefore, an orderly plan is required to separate the ions and correctly identify each.
In part A of this experiment, you will follow an orderly plan for identifying the ions of the
halogen family. You will then attempt to identify an unknown substance that contains one
of the ions from the halogen family.
Ch6 HW#4
Ch6 HW#4
1. Which of these elements has a greater electronegativity?
a. Na or K
b. Mg or P
c. Fr or F
2. Use the Electronegativity Table to decide whether each pair of atoms will form
an ionic or covalent bond:
a. Na and Br
b. Ca and Cl
c. C and O
d. Al and O
e. Mg and S
Ch9.1 Naming Ions
Elements classified as metals, nonmetals, and metalloids.
Ch9.1 Naming Ions
Elements classified as metals, nonmetals, and metalloids.
Metals form Cations - give up electrons so that the valence shell is
empty. (Positively Charged)
Lose as many electrons as their group #.
Exs:
K 
Be 
B 
These ions are named the same as their neutral atom.
Ch9.1 Naming Ions
Elements classified as metals, nonmetals, and metalloids.
Metals form Cations - give up electrons so that the valence shell is
empty. (Positively Charged)
Lose as many electrons as their group #.
Exs:
K  K1+ potassium ion
Be  Be2+ beryllium ion
B  B3+ boron ion
These ions are named the same as their neutral atom.
Nonmetals form Anions - gain electrons so that its valence shell is
full. (Negatively Charged)
Gain electrons based on: (group #) + (electrons) = 8
Exs:
N 
Se 
I 
Elements classified as metals, nonmetals, and metalloids.
Metals form Cations - give up electrons so that the valence shell is
empty. (Positively Charged)
Lose as many electrons as their group #.
Exs:
K  K1+ potassium ion
Be  Be2+ beryllium ion
B  B3+ boron ion
These ions are named the same as their neutral atom.
Nonmetals form Anions - gain electrons so that its valence shell is
full. (Negatively Charged)
Gain electrons based on: (group #) + (electrons) = 8
Exs:
N  N3– nitride ion
Se  Se2– selenide ion
I  I1– iodide ion
These ions are named by taking the name of the neutral atom and
and changing the ending to “-ide”
Elements classified as metals, nonmetals, and metalloids.
Metals form Cations - give up electrons so that the valence shell is
empty. (Positively Charged)
Lose as many electrons as their group #.
Exs:
K 
Be 
B 
K1+ potassium ion
Be2+ beryllium ion
B3+ boron ion
These ions are named the same as their neutral atom.
Nonmetals form Anions - gain electrons so that its valence shell is
full. (Negatively Charged)
Gain electrons based on: (group #) + (electrons) = 8
Exs:
N  N3– nitride ion
Se  Se2– selenide ion
I  I1– iodide ion
These ions are named by taking the name of the neutral atom and
and changing the ending to “-ide”
Group 14 and 18 elements generally do not form ions.
Transition metals can form multiple ions, depending on how many
electrons lost. Refer to your oxidation sheet for common
transition metal oxidation numbers.
Most polyatomic ions end in “-ate” or “-ite”
The most common from a group is named with “-ate”
The ion with one less oxygen gets “-ite”
One less
oxygen
Most
common
-ite:
-ate:
SO4–2
NO3–
ClO3–
Most polyatomic ions end in “-ate” or “-ite”
The most common from a group is named with “-ate”
The ion with one less oxygen gets “-ite”
Two less oxygens gets “hypo” prefix and “-ite” ending
One more oxygen than most common gets “per” prefix and “-ate” ending
Two less
One less
Most
One more
oxygens
oxygen
common
oxygen
hypo(root)ite:
-ite:
SO3–2
NO2–
ClO2–
-ate:
SO4–2
NO3–
ClO3–
per(root)ate
Most polyatomic ions end in “-ate” or “-ite”
The most common from a group is named with “-ate”
The ion with one less oxygen gets “-ite”
Two less oxygens gets “hypo” prefix and “-ite” ending
One more oxygen than most common gets “per” prefix and “-ate” ending
Two less
One less
Most
One more
oxygens
oxygen
common
oxygen
hypo(root)ite:
ClO–
-ite:
SO3–2
NO2–
ClO2–
-ate:
SO4–2
NO3–
ClO3–
per(root)ate
ClO4–
Other endings include “-ium” (like ammonium, NH31+ )
and the occasional “-ide” (like cyanide, CN– and hydroxide, OH–)
Ch9 HW#1
Ch9 HW#1
1. Identify the following elements as metals, metalloids, or nonmetals.
a. gold
b. silicon
c. manganese
d. sulfur
e. barium
2. Describe two ways in which an ion forms from an atom.
3. Complete this table.
a.
b.
c.
d.
e.
f.
g.
Symbol
Al
S
Change in electrons
3 electrons lost
Br
1 electron gained
Formula of Ion
Name of Ion
S–2
Ca+2
P–3
Xe
K
4. The most common polyatomic ion of bromine is bromate, BrO3–.
Name and write the formulas of the bromine ions with 4 oxygens and also 1 oxygen.
Ch9.2 Chemical Formulas and Naming
Some substances have common names
Ex: H2O:
Just about all others we will name systematically:
CO2:
NaCl:
Why the difference?
Ionic compounds
- as a general rule: metal and a nonmetal
- cation listed before the anion, anion changed to “-ide.”
Ex1) Li and F
Ca and S
Na and I
Ex1) Li and F  LiF  Lithium Fluoride
Ca and S  CaS  Calcium Sulfide
Na and I  NaI  Sodium Iodide
(Tomorrow we can learn how to write formulas.)
- the total charge of the compound is zero.
The ions always combine in ratios
to cancel each other out.
molecular compounds
- usually consist on nonmetals bonded together
- share electrons in a covalent bond
- usually neutral in charge
(polyatomic ions are covalently bonded elements with an overall charge.)
- still end in “-ide”
- naming based on prefixes:
Ex2) SO2
N2O4
BF3
What would be another name for water?
Prefix
Number
Mono-*
1
Di-
2
Tri-
3
Tetra-
4
Penta-
5
Hexa-
6
Hepta-
7
Octa-
8
Nona-
9
Deca-
10
Here are some more properties of compounds formed
from ionic and covalent bonds:
Ch9 HW#2
Ch9 HW#2
5. If ionic compounds are composed of charged particles (ions),
why isn’t every ionic compound either positively or negatively charged?
6. The melting point of a compound is 1240°C.
Do you predict this compound to be an ionic or molecular compound?
7. Identify the number and kinds of atoms present in a molecule
of each compound:
a. ascorbic acid (vitamin C) C6H806
b. monosodium glutamate (MSG) C5H8O4NNa
c. sucrose (table sugar) C12H22O11
d. trinitrotoluene (TNT) C7H5N3O6
e. ammonium nitrate (fertilizer) NH4NO3
8. Would you expect the following pairs of atoms to combine chemically
to give an ionic or a molecular compound?
a. Li and S
b. O and S
c. Al and O
d. F and Cl
e. I and K
f . H and N
9. Name these molecular compounds:
a.
b.
c.
d.
e.
CS2
Cl2O7
N2O5
CCl4
CrCl3
10. Name these ionic compounds:
a.
b.
c.
d.
e.
KBr
CaCl2
MgO
LiCl
Fe2O3
Ch9.3 Chemical Formulas
Binary compounds – made up of only two kinds of atoms.
Ternary compounds – made up of more than two kinds of atoms.
Ex1) Write chemical formulas for the following:
sodium fluoride
calcium sulfide
magnesium chloride
lithium nitride
aluminum oxide
Ex2) Write chemical formulas for the following:
sodium nitride
sodium nitrate
calcium carbide
calcium carbonate
magnesium sulfide
magnesium sulfite
magnesium sulfate
Ch9 HW#3
There will be quizzes on formula writing ALL year!
Might as well learn it now! Do your HW!
Ch9 HW#3
11. What is the net ionic charge of every ionic compound?
12. Write formulas for compounds composed of these pairs of ions:
a. Na+, O-2
b. Sn+4, S-2
c. K+, Cld. Mg+2, N-3
13. Write the formulas for these compounds.
a. potasium chloride
b. potasium hypochlorite
c. potasium chlorite
d. potasium chlorate
e. potasium perchlorate
14. Write names for these compounds.
a. Al(OH)3
b. NaI
c. Na2CrO4
Ch9.4 More Naming
Transition metals can have more than one oxidation number.
This is denoted with Roman Numerals
Ex1) Write the chemical formula and name for:
Sn4+ and oxygen
Fe2+ and oxygen
Fe3+ and oxygen
Ex2) Write the chemical formula for:
lead (I) carbonate
copper (II) sulfate
silver nitrate
Ch9 HW#4
Ch9 HW#4
15. How do you determine the charge of a transition metal cation
from the formula of an ionic compound?
16. Write the names for these binary ionic compounds:
a. AlF3
b. FeS
c. CaO
d. Cu2Se
e. BF3
17. Write the formulas for:
a. calcium carbonate
b. lithium hypochlorite
c. potassium phosphate
d. tin (IV) dichromate
Ch4 Rev WS on Formulas and Naming
(30 questions, we do 10, you do 2 for HW)
Ch9.5 Naming Acids and Bases (Discuss these in more detail later)
Acids: A compound that contains one or more hydrogen atoms and
produces hydrogen ions (H+) when dissolved in water. (The names
of compounds can change when acids are present).
The name of an acidic compound depends on the type of anion
it is reacting with...
For “–ide” anions, the prefix “hydro-” comes before the root
word in the anion and the suffix “-ic” comes after.
Ex: HCl
For “-ite” anions, just the suffix “-ous” comes after the root word.
Ex: HClO2
For “-ate” anions, just add the suffix “-ic” after the root word.
Ex: HClO3
The word “acid” is then added after each of these names.
Bases: An ionic compound that produces hydroxide ions (OH-) when
dissolved in water.
Naming is similar to any ionic compound:
Ex: NaOH
Chemical formulas of Acids and Bases:
calcium hydroxide
nitric acid
nitrous acid
Ch9 HW#5
Ch9 HW#5
18. The formula of a common acid will always contain what element?
19. Name these acids and bases. Please feel free to look on your
oxidation sheet to find the name of the polyatomic ion.
a. HF
b. H2S
c. H2SO4
d. H2SO3
e. KOH
20. Write the chemical formula of the following acids and bases.
a. Hydrobromic acid
b. Carbonic acid
c. Phosphoric acid
e. Aluminum hydroxide
d. Magnesium hydroxide
Ch6 Review
1. Between lithium and potassium, which element has
a. the largest atomic radius
__________
b. the highest ionization energy
__________
c. the highest electronegativity
__________
2. Between sodium and aluminum, which element has
a. the largest atomic radius
__________
b. the highest ionization energy
__________
c. the highest electronegativity
__________
3. Between sulfur, arsenic and selenium, which element has
a. the largest atomic radius
__________
b. the highest ionization energy
__________
c. the highest electronegativity
__________
4. Which has the larger size:
a. calcium or its ion
__________
b. aluminum or its ion
__________
c. chlorine or its ion
__________
5. “Periodic” means that something
repeats or goes through cycles.
From the graph do you see any
cyclical nature to the points?
6. List the elements that represent
the peaks in your graph.
How are these elements related
on the periodic table?
7. List the elements that represent
the lower peaks in your graph.
How are these elements related
on the periodic table?
8. List the elements that occur just before each peak
on your graph. How are these elements related on the periodic table?
9. List the elements that occur just after each lower peak on your graph.
How are these elements related on the periodic table?
10. The following graph shows the successive ionization energies for period 3
on the periodic table. Referring to the diagram of aluminum,
explain the first four ionization energies.
13p
14n
Ch9 Rev
1. Would you expect the following compounds to be ionic or molecular?
a. CO
b. Na2S
2. Write the formulas for these compounds.
a. sodium iodide
b. stannous fluoride (Sn2+)
c. potassium sulfide
d. calcium iodide
c. C2H6
d. N2O3
3. Name these binary ionic compounds.
a. ZnS
b. KCl
c. BaO
d. CuBr2
4. Write formulas for the following binary molecular compounds.
a. nitrogen trifluoride
b. disulfur dichloride
c. dinitrogen tetroxide
d. phosphorus pentachloride
5. When are parentheses needed in a chemical formula?
Use barium hydroxide to explain.
6. The formula of a common acid will always contain what element?
7. Name these compounds.
a. CaCO3
b. KMnO4
c. PbCrO4
d. Mg3P2
e. (NH4)2SO4
f. NaNO2
8. Write the formulas for these compounds.
a. tin(II) hydroxide
b. barium fluoride
c. calcium sulfide
d. sodium phosphate