The Periodic Table Learning Targets (Your exam at the end of Unit 2 ...3+Packet+Honor… · Chemistry HP Unit 3 – The Periodic Table Learning Targets (Your exam at the end of Unit

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Chemistry HP Unit 3 – The Periodic Table

Learning Targets (Your exam at the end of Unit 2 will assess the following:)

3. The Periodic Table

3-1. Discuss the development of the periodic table by Mendeleev.

3-2. Locate and state important properties of main chemical families including the alkali metals, alkaline earth

metals, transition metals, halogens, noble gases, lanthanides, actinides, and hydrogen.

3-3. Define atomic radius explain periodic trends in this property as they relate to atomic structure.

3-4. Define ionization energy and explain periodic trends in this property as they relate to atomic structure. List

elements that are exceptions to the general trend and use orbital notation to show why they are exceptions.

3-5. Define electronegativity and explain periodic trends in this property as they relate to atomic structure.

3-6. Define ionic radius and relate the size of an anion to a neutral atom of the same element and a cation to a

neutral atom of the same element.

3-7. Draw electron dot diagrams for atoms, showing the correct number of valence electrons.

3-8. Draw Lewis structures from chemical formulas.

3-9. Assign bond orders for a molecule from the Lewis structure

3-10. Calculate the total number of valence electrons in a polyatomic ion.

3-11. Draw Lewis structures for polyatomic ions.

3-12. Assign formal charges to atoms in polyatomic ions.

3-13. Draw resonance structures for polyatomic ions.

3-14. Classify bonds as ionic or covalent and as polar or non-polar using electronegativity values.

3-15. Assign shapes to molecules using VSEPR Theory and draw the VSEPR diagrams for a molecule.

3-16. Classify molecules as polar or non-polar using shape.

3-17. Compare miscible and immiscible, by definition and with example, and determine if two substances will be

miscible or immiscible based on polarity

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3-1. Discuss the development of the periodic table by Mendeleev.

Mendeleev’s Periodic Table

Dmitri Mendeleev was a Russian chemistry professor who was looking for a way to ____________________ the

____________________ ________________________ to aid his students.

He listed the elements in ______________________ according to their ____________________ _____________________, and in

_____________________when _____________________ of the elements began to _____________________.

Mendeleev left ____________________ in his table for elements that had yet to be discovered.

When the elements _____________________ and _______________________ were later discovered, they fit perfectly into the gaps

Mendeleev had left for them.

Mendeleev also believed that the placement of elements ___________________ and _____________________ should be switched,

to correspond with their chemical properties.

When the Modern Periodic Table was later organized according to _______________________ ___________________ and not

_______________________ ___________________, the elements ________________________ and ____________________ lined

up correctly according to their chemical properties.

Mendeleev Periodic Table CLOZE:

Mendeleev organized elements by _______________ ________________, NOT__________________ __________________

(__________________ __________________ had not been discovered yet), AND organized them by similar

__________________ . He left _______________ for ________________________ elements which predicted the

________________________ of those elements, which were discovered later.

The Modern Periodic Table The Modern Periodic Table organizes elements into rows, called _____________________, according to ___________________ ______________, and NOT ___________________ ______________. Each period corresponds to the ________________ __________________, or ________________ __________________ ________________, _____________. The elements are also organized into ________________________ according to _______________________ _____________________. The columns are called _____________________. Each column, corresponds to a different ________________ of elements

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3-2. Locate and state important properties of main chemical families including the alkali metals, alkaline earth metals,

transition metals, halogens, noble gases, lanthanides, actinides, and hydrogen.

Chemical Groups/Families

Elements in the same chemical family share similar properties. Some chemical families have been given names.

Group Number and Name

Elements Ion(s) Formed Properties

Hydrogen

I or 1 Alkali Metals

II or 2 Alkaline Earth Metals

3-12 Transition Metals

VII or 17 Halogens

VIII or 18 Nobel Gases

Lanthanides (57-70) and Actinides (89-102)

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WS #1 (Learning Target 3-2: Locate and state important properties of main chemical families including the alkali metals,

alkaline earth metals, transition metals, halogens, noble gases, lanthanides, actinides, and hydrogen.)

Periodic Table Puzzle

Fictitious symbols are used for the first 18 elements in the periodic table. Use the clues below to write the fictitious symbol in the appropriate spot on the periodic table provided. Symbols for real elements do not represent those elements. HINT: You do not have to complete each clue in order.

Clue 1 U and J are alkali metals. J has more energy levels.

Clue 2 T has 4 valence electrons on the 3rd energy level.

Clue 3 M is a metal in period 3 with 2 valence electrons.

Clue 4 X has one proton in its nucleus.

Clue 5 Q has 2 energy levels, is a nonmetal, and is a solid at room temperature.

Clue 6 L is a noble gas that doesn’t have 8 valence electrons.

Clue 7 Z and Y are members of the nitrogen family. Y is a gas at room temperature.

Clue 8 D has an ending electron distribution of s2p5. R has an ending electron distribution of s2.

Clue 9 G has 6 valence electrons.

Clue 10 V and W have full outer energy levels. V has 3 energy levels.

Clue 11 A atoms have 3 valence electrons and E atoms have 6 valence electrons. Both are in the second period.

Clue 12 K has one fewer total electrons than V.

Clue 13 I has 3 valence electrons on the third energy level.

Answers:

X L

U-R-A-Q-Y-E-D-W

J-M-I-T-Z-G-K-V

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3-3. Define atomic radius explain periodic trends in this property as they relate to atomic structure.

Periodic Trends

Some chemical properties of elements vary predictably down a group and across a row on the periodic table. These tendencies are called periodic trends and are based on atomic structure. Periodic trends are observed in the properties of atomic radius, ionization energy, and electronegativity.

Trend Atomic Radius

Definition Atomic radius is the ________ of an atom. The atomic radius is measured as ________________ from the nucleus to the ________________ ________________ _________________. Atomic radius is usually measure in ________________ (1 pm = _________ m). Atoms range in size from __________ (Helium) to ___________ (Francium).

Down a Group Atomic Radius _________________ down a group. Atoms become larger down a group of the

periodic table. Each new element down a group adds a ____________ ______________

______________ . Since electrons are located in _____________ ________________

________________, they are _______________ from the nucleus, making atoms ______________.

Across a Period Atomic Radius ______________ across a row. Atoms become smaller across a row of the periodic

table. Across a row, _________new energy levels are added, but the number of _____________in

the _______________ increases. Therefore, the positive charge on the nucleus becomes

_____________ and electron shells are ________________ _____________ __________________

to the nucleus, making atoms ___________________.

Example Which atom has a larger atomic radius? Why?

(1) Na or K?

(2) C or O?

Atomic Radius CLOZE Atomic radius is defined as the ______________ from the nucleus to the outermost electron in an atom. It is usually measured in the unit _______________. As you go down a _____________ on the Periodic Table, ionic radius ________________, because, with each new period, an additional _______________ ____________ is added, making the atomic radius ____________. Therefore, we would expect the atomic radius of Br to be ___________ than that of Cl. As you go across a period on the Periodic Table, atomic radius ______________. This is because the number of protons _________________ as you go left to right, while the energy level _____________ _________ ____________. The increased number of _______________ creates a greater ______________ force between the nucleus and its electrons, bringing the outermost electrons _______________ in, and making the atomic radius ______________. Therefore, we would expect the atomic radius of Br to be _____________ than that of Se.

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3-4. Define ionization energy and explain periodic trends in this property as they relate to atomic structure. List elements

that are exceptions to the general trend and use orbital notation to show why they are exceptions.

Trend Ionization Energy

Definition Ionization energy is the ___________ required to remove the __________________ ___________________ from an atom. For example, the ionization for an atom of lithium: Li + IE → Li+ + e- The ionization for an atom of fluorine: F + IE → F+ + e- Generally, the closer an electron is to the nucleus, the ________________ it is to remove and the

________________ the ionization energy will be.

Down a Group Ionization Energy __________________ down a group.

Each new element down a group adds a ___________ _____________________ ______________,

so the outermost electrons are located in shells ___________________ from the nucleus, making

the outermost electrons _________________ to remove.

Across a Period Ionization Energy ___________________ across a row. Across a row, ________ ___________

_________________ _______________ are added, but the number of protons in the nucleus

______________________. Therefore, the electrons are pulled __________________ to the

nucleus and are _______________________ to remove, __________________ the amount of

_______________________ to remove the electrons, __________________ the ionization energy.

Example Which atom has a higher IE? Why?

(1) Na or K?

(2) C or O?

Exceptions Ionization Energy and Ion Stability

Ionization energy can also depend on the __________________ of the ____________

____________ by losing the electron. In general, ions are most stable when the highest energy

orbitals are _________________ or _____________- _____________. The ___________________

______________________ of an atom must be considered to explain some trends.

Example Which atom has a higher IE? Why?

(1) Al or Mg?

We would expect Al to have a ____________ ionization energy than Mg because its outermost

electrons are in the _______________ ______________, and Mg has ____________ protons

than Al. However, Mg actually has a ________________ ionization energy than Al because:

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When Mg loses an electron, Mg → Mg+ + e-, its electron configuration changes from 1s2 2s2 2p6

3s2 to 1s2 2s2 2p6 3s1 + e-. In doing so, it goes from a ____________ 3s orbital to an

_____________ 3s orbital. On the other hand, when Al loses an electron, Al → Al+ + e-, its

electron configuration changes from 1s2 2s2 2p6 3s2 3p1 to 1s2 2s2 2p6 3s2 + e-. In doing so, Al

goes from an _______________ 3p orbital to a _______________ 3s orbital. Therefore, the

ionization energy of Al is _____________ than the ionization energy of Mg.

(2) S or P?

Ionization Energy CLOZE Ionization Energy is defined as the ______________ required to remove the outermost electron from an atom.

As you go down a _____________ on the Periodic Table, ionization energy ________________, because, with each

new period, an additional _______________ ____________ is added, bringing the outermost electron

_____________ _____________ the nucleus. Therefore, we would expect the ionization energy of Br to be

___________ than that of Cl.

As you go across a period on the Periodic Table, ionization energy ______________. This is because the number of

protons _________________ as you go left to right, while the energy level _____________ _________ ____________.

The increased number of _______________ creates a greater ______________ force between the nucleus and its

electrons, making it ______________ to remove an electron. Therefore, we would expect the ionization energy of Br

to be _____________ than that of Se.

An exception to this trend is when the ion that is formed by removing the outermost electron is more stable than the

neutral atom. When this is the case, ionization energy across a period may not always ______________ from one

element to the next. Boron and Beryllium are examples of this. One would expect the ionization energy of boron to

be _____________ than that of beryllium. But, since removing one electron from boron results in a filled ________

sublevel, whereas removing one electron from beryllium removes an electron from a filled ___________ sublevel, the

ionization energy of boron is _______________ than that of beryllium.

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3-5. Define electronegativity and explain periodic trends in this property as they relate to atomic structure.

Trend Electronegativity

Definition Electronegativity is the measure of the ability of an atom in a chemical compound to

________________ the electrons of the other atom in a _______________.

The closer the electrons of another atom in a bond can be to the __________________, the

___________________ the electronegativity.

In general, _______________ atoms have a ________________ electronegativity and

________________ atoms have a _______________electronegativity.

Electronegativity values are measured on a scale from ______ to ______.

Down a Group Electronegativity _________________ down a group. Down a group, atoms become

_________________, so the electrons of another atom in a bond would be ______________

________________ from the nucleus and electronegativity is ___________________.

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Across a Period Electronegativity _________________ across a row. Across a row, atoms become

______________, so electrons of another atom in a bond can be _______________ to the nucleus

and electronegativity is ______________________.

Example Which atom has a higher EN? Why?

(1) Na or K?

(2) C or O?

Electronegativity CLOZE Electronegativity is defined as the ability of an atom in a chemical bond to attract the other atom’s ________________. As you go down a _____________ on the Periodic Table, electronegativity ________________, because, with each

new period, an additional _______________ ____________ is added, increasing the size of the atom. Since the larger

the atom, the ________________ the nucleus is from its outermost electrons, and the _________________ the

nucleus is from the electrons of the other atom in a chemical bond, it becomes more difficult to attract the other

atom’s electrons toward itself. Therefore, we would expect the electronegativity of Br to be ___________ than that

of Cl.

As you go across a period on the Periodic Table, electronegativity ______________. This is because the number of

protons _________________ as you go left to right, while the energy level _____________ _________ ____________.

The increased number of _______________ creates a greater ______________ force between the nucleus and its

electrons, and between the nucleus and the electrons of the other atom in a chemical bond, making it

______________ to attract an electron. Therefore, we would expect the electronegativity of Br to be _____________

than that of Se.

In general, noble gases have an electronegativity of _________________, because they usually do not form chemical

bonds.

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3-6. Define ionic radius and relate the size of an anion to a neutral atom of the same element and a cation to a neutral

atom of the same element.

Trend Ionic Radius

Definition Ionic radius is the _______________ of the ion of an element compared to the neutral atom of the

element.

Cations Positive Ions/Cations Atoms become positively charged ions by ______________ electrons. Since

electrons are lost, a positive ion is _______________ than the neutral atom.

Anions Negative Ions/Anions Atoms become negatively charged ions by ________________ electrons.

Since electrons are gained, a negative ion is ___________________ than the neutral atom

Example Which has a larger radius, the neutral atom or the ion? Why?

(1) Ca or Ca2+?

(2) S or S2-?

Ionic Radius CLOZE Ionic radius is defined as the __________________ from the nucleus to the outermost electron in an ion.

When an atom _________________ an electron, becoming an anion, its ionic radius ___________________.

Therefore, we would expect the radius of Br- to be ________________ than the atomic radius of neutral Br.

When an atom _________________ an electron, becoming a cation, its ionic radius ____________________.

Therefore, we would expect the radius of Na+ to be ________________ than the atomic radius of neutral Na.

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WS #2 (Learning Targets 3-3, 3-4, 3-5, 3-6)

Periodic Trends

(1) List each set of atoms from smallest to largest radius.

(a) Mg, Be, Ca (b) C, B, N

(c) As, P, N (d) Sc, Ti, V

(2) Which of the elements listed has the highest ionization energy?

(a) K, Li, Cs (b) C, F, Be

(c) Na, Mg, Cl (d) Ra, Ca, Be

(3) Which of the elements listed has the highest electronegativity?

(a) Be, B, N (b) Se, O, S

(c) I, Br, At (d) Na, Al, P

(4) Complete the following table:

Give Electron Configuration (a) Si or Cl?

Si:

Cl:

(b) F or Br?

F:

Br:

Which element has a larger atomic radius? Explain.

Which element has a higher ionization energy? Explain.

Which element has a higher electronegativity? Explain.

(5) Which has a larger radius, the neutral atom or the ion? Explain.

(a) Br or Br˗ (b) Cs or Cs+

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Atoms react with each other and form bonds in order to ________________ ________________ ________________

________________ ________________ ________________with electrons. This is why noble gases generally do not bond with other

elements; they already have full outer shells.

The types of bonds they form depend upon the difference in electronegativity between the two atoms forming the bond. Recall that

electronegativity is a measure of the ability of an atom in a chemical compound to ________________ ________________

________________ ________________ ________________ ________________ ________________ in a chemical bond. A bond in

this sense then, is a tug of war between electrons. The greater the electronegativity difference between the two atoms forming a

bond, the more the electrons will ________________ ________________ around the ________________ ________________atom.

As an example, let’s consider the bond sodium and chlorine form when making sodium chloride, also known as table salt. Both

sodium and chlorine are in period ________________ of the periodic table. Sodium is in the first group, the ________________

________________, whereas chlorine is a ________________, in the second to last group.

According to the periodic trends, we would expect chlorine to have a ________________ electronegativity than sodium, because

chlorine has a ________________ nuclear charge (it has more ________________ in its nucleus), while both sodium’s and chlorine’s

outermost electrons are ________________ distances from the nucleus, being in the same ________________.

We can confirm this by comparing the electronegativity values of sodium and chlorine.

From the chart, you can see that the electronegativity of chlorine is ________________, and the electronegativity of sodium is

________________. That’s an electronegativity difference of ________________- ________________= ________________. This

electronegativity difference is so ________________, that chlorine literally ________________ sodium’s valence electron, forming

the ________________ ________________ ________________, leaving sodium as the ion ________________. The bond that

sodium and chlorine form is therefore called an ________________ ________________, because the electronegativity difference

between the elements is so great that the valence electrons are literally ________________ from the ________________

electronegative atom to the ________________ electronegative atom, turning each atom into an ________________.

An ionic bond can therefore be defined as a bond between ________________ ________________, caused by an electronegativity

difference of ________________ or greater. What holds these ions together in an ionic bond is the ________________

________________ between the cation (________________ ion) and the anion (________________ ion). The rule of thumb is that

an ionic bond forms when a ________________ bonds with a ________________. This makes sense when you consider that metals

are on the ________________ side of the periodic table and nonmetals are on the ________________ side. Nonmetals have a

________________ electronegativity then metals, because they generally have a ________________ nuclear charge.

Atoms whose electronegativity difference lies between ________________ and ________________ are called ________________

bonds. The term covalent is made up of the prefix “co-” which is a ________________ ________________ , and “-valent” which

refers to the ________________ ________________ . Covalent bonds are therefore a ________________ of electrons. Neither

of the two atoms in a covalent bond have sufficiently more electronegativity than the other to literally take the other’s valence

electrons. Rather, in a covalent bond, the valence electrons are ________________ .

An example of a covalent bond is the bond between oxygen and hydrogen in the molecule water. Hydrogen has an electronegativity

of ________________ and oxygen has an electronegativity of ________________. The electronegativity difference between these

two atoms is ________________- ________________= ________________. Since ________________is ________________ than

1.7, oxygen and hydrogen form a ________________ bond.

Because oxygen has a ________________ electronegativity than hydrogen, the electrons will ________________

________________more around the oxygen, giving the oxygen an excess ________________ charge, and leaving the hydrogen with

an excess ________________ charge. But, because the electronegativity difference is ________________ high enough, the

electrons will ________________ completely ________________ hydrogen to join oxygen.

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A covalent bond can therefore be defined as a bond between two atoms in which the electronegativity difference of is less than or

equal to ________________. The rule of thumb is that a covalent bond forms when ________________ bond with

________________ ________________. This makes sense when you consider that all the nonmetals are on the ________________

side of the periodic table, so their electronegativity values are ________________ to each other.

Covalent compounds can be further classified as ________________ ________________ or ________________ ________________.

Polar comes from the word ________________, describing the existence of positive and negative regions, or

________________within a molecule. In the last example comparing the electronegativities of hydrogen and oxygen, we saw that

the electronegativity difference between hydrogen and oxygen was ________________, which is less than the

________________difference required for ionic compounds. However, since there ________________ a difference in

electronegativity, the electrons in the compound will be attracted to the ________________ atom more than the

________________ atom. This will leave the oxygen atom with a ________________ ________________ charge, also called a

________________ ________________, and the hydrogen atom with a ________________ ________________charge because of its

________________ of electron density, also called a ________________ ________________.

We draw these poles using the following symbols:

The bond between oxygen and hydrogen in water is referred to as a ________________ ________________ ________________,

because, even though the bond is covalent, it does contains positive and negative ________________.

If the electronegativity difference between two atoms in a covalent compound is less than or equal to ________________, we refer

to the bond as ________________ ________________. If the electronegativity difference is between ________________and

________________, we refer to the bond as ________________ ________________.

The following chart summarizes these electronegativity differences:

Electronegativity Difference

Covalent Ionic

Nonpolar Covalent Polar Covalent

0 to 0.4 0.5 to 1.7 1.8 and up

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To Read Before the Next Lesson:

The type of bond shared between elements in a compound determines some of the properties of that compound. Ionic compounds, formed when ions bond together, exist as crystal structures with repeating cation-anion arrangements, such as the example drawn below of NaCl.

The large green spheres are the sodium ions, and the smaller blue spheres are the chloride ions. They assemble into this crystal lattice structure to maximize electrostatic attraction between the cations and anions, and to minimize the electrostatic repulsion between like charges. Here is a 3D image of the same structure:

Since a sodium ion is bonded within an entire crystal of repeating positive and negative ions, we don’t refer to NaCl as a molecule, but rather a formula unit. NaCl is the formula unit of the repeating Na-Cl-Na-Cl-Na-Cl crystalline structure. Because of this strong electrostatic attraction between the positive and negative ions, the melting points of ionic compounds are high. It takes a lot of energy to loosen the hold these ions have on each other. In addition, ionic compounds are brittle. And, they act as electrolytes, conducting electricity in solution.

In summary, the properties of ionic compounds are as follows: 1. They exist as repeating units within a crystalline structure. 2. High melting points. 3. Brittle 4. Strong electrolytes (conduct electricity when in solution). Covalent compounds are the result of nonmetals bonding with each other and sharing their electrons in such a way that every atom in the bond has a full set of valence electrons. Covalent compounds are referred to as molecules, not formula units. In general, covalent compounds have: 1. Lower melting and boiling points. 2. Are non-electrolytes (They do NOT conduct electricity). 3. Exist as mostly liquids or gases at room temperature.

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WS #3 (Learning Target 3-14. Classify bonds as ionic or covalent and as polar or non-polar using electronegativity values.)

Directions: Classify the following bonds as ionic, polar or nonpolar covalent. Draw the bond. If the bond is polar

covalent, indicate the distribution of charge using the arrow and delta symbols.

Compound Electronegativity Difference Type of Bond Draw Bond

(a) Br2

(b) MgO

(c) LiF

(d) SeO

(e) ICl

(f) BrCl

(g) CO

(h) NaCl

(i) HCl

(j) O2

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Answers.

Compound Electronegativity Difference Type of Bond Draw Bond

(a) Br2 0

(b) MgO

(c) LiF

(d) SeO

(e) ICl

(f) BrCl

(g) CO

(h) NaCl

(i) HCl

(j) O2

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3-7. Draw electron dot diagrams for atoms, showing the correct number of valence electrons.

We represent both ionic and covalent bonds using a notation known as Lewis Dot Structures, which are a way of representing the

valence electrons around an atom.

How to write Lewis Dot Structures (electron dot diagrams)

1. Write the chemical symbol of the element.

2. Put one dot for each valence electron.

3. Don’t pair electrons until you have to.

Element Number of Valence Electrons

Lewis Dot Structure Ion Number of Valence Electrons of Ion

Lewis Dot Structure of Ion

N

Cl

Al

Mg

He

Na

Lewis Dot Structures of Ionic Compounds

Draw the Lewis Dot Structures of the Ions making up the ionic compound next to each other.

Ionic Compound Lewis Dot Structure of Cation

Lewis Dot Structure of Anion

Lewis Dot Structure of Ionic Compound

NaCl

MgO

BeS

MgCl2

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WS #4 (Learning Target 3.7: Draw electron dot diagrams for atoms, showing the correct number of valence electrons. )

Directions: Draw the Lewis Dot Structures for the following atoms, ions and ionic compounds.

Atom, Ion or Ionic Compound Lewis Dot Structure

1) Ca2+

2) I-

3) Ge

4) As

5) Rb+

6) RbI

7) AlCl3

8) CaO

9) LiBr

10) Na2S

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To Draw Lewis Dot Structures for Covalent Compounds, follow these 5 steps:

1.

2.

3.

4.

5.

Example: Draw the Lewis Dot Structure for CF4

What is the bond order for each C-F bond? ______

Example: Draw the Lewis Dot Structure for H2O

What is the bond order for each O-H bond? ______

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Example: Draw the Lewis Dot Structure for CO2

What is the bond order for each C-O bond? ______

Example: Draw the Lewis Dot Structure for N2

What is the bond order for the N-N bond? ______

Exceptions:

Incomplete Octet Expanded Valence

Examples: Examples:

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WS #5 (Learning Target 3-8, 3-9)

Directions: Draw Lewis structures for the following compounds.

(a) I2 (b) CF4

(c) NCl3 (d) SCl2

(e) SiTe2 (f) PF3

(g) GeBr4 (h) TeCl2

(i) CSe2 (j) H2

(k) S2 (l) P2

(m) C2 (n) C2H6

(o) C2H4 (p) C2H2

(q) HSCN (r) BBr3

(s) PI5 (t) SAt6

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3-10. Calculate the total number of valence electrons in a polyatomic ion.

3-11. Draw Lewis structures for polyatomic ions.

3-12. Assign formal charges to atoms in polyatomic ions.

3-13. Draw resonance structures for polyatomic ions.

A polyatomic ion is a group of _____________ ___________ ___________ ___________– ___________-- acting together as

___________ ___________ ___________.

Examples:

When drawing the Lewis structure for a polyatomic ion, _____________ _____________ _____________for each _____________

charge and _____________ _____________ _____________for each _____________ charge.

Example: Draw the Lewis Dot Structure for CN-.

What is the bond order for the C-N bond? ______

Formal Charge

The charge on each atom in the compound can be determined by calculating the _____________ _____________. The formal

charges for all atoms must add up to the _____________ charge for the polyatomic ion.

The formal charge on each atom in a polyatomic ion can be calculated as follows:

Formal Charge = number of valence electrons – number of unshared electrons – ½ number of shared (bonding) electrons

Formal Charge of CN- # Valence Electrons # Non-bonding Electrons ½ # Bonding Electrons Formal Charge

Carbon

Nitrogen

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Example: Draw the Lewis Dot Structure for NO3-.

What is the bond order for each N-O bond? ______

Formal Charge of NO3- # Valence Electrons # Non-bonding Electrons ½ # Bonding Electrons Formal Charge

Nitrogen

Oxygen (1,2)

Oxygen (3)

Resonance Structures

Resonance Structures = When 2 or more ______________ _____________ _____________ are necessary to describe the bonding in

a molecule or ion.

Example: Resonance Structures of NO3-.

What is the bond order for each N-O bond? ______

When do you draw resonance structures?

1. When you have more than one possible Lewis Dot Structure.

2. Common in Polyatomic Ions containing oxygen.

Bond Order of Resonance Structures = Total Number of Bonds ÷ Total Number of Bonding Locations.

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WS #6 (Learning Targets 3-10, 3-11, 3-12, 3-13)

(1) Determine the total number of valence electrons in each of the following polyatomic ions. Draw Lewis structures and

calculate the formal charge on each atom.

(a) OCl– (b) OH–

(c) H3O+ (d) CO3 2–

(e) NO2 – (f) PO3 3–

(g) CN– (h) SO4 2–

(2) Draw all possible resonance structures for the polyatomic ion. Calculate the formal charge on each atom. Determine

the bond order.

(a) SiO3 2– (b) PO4 3–

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Answers.

(1) Determine the total number of valence electrons in each of the following polyatomic ions. Draw Lewis structures and

calculate the formal charge on each atom.

(a) OCl– (b) OH–

(c) H3O+ (d) CO3 2–

(e) NO2 – (f) PO3 3–

(g) CN– (h) SO4 2–

(2) Draw all possible resonance structures for the polyatomic ion. Calculate the formal charge on each atom. Determine

the bond order.

(a) SiO3 2– (b) PO4 3–

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3-15. Assign shapes to molecules using VSEPR Theory and draw the VSEPR diagrams for a molecule.

3-16. Classify molecules as polar or non-polar using shape.

3-17. Compare miscible and immiscible, by definition and with example, and determine if two substances will be miscible

or immiscible based on polarity

VSEPR Theory = Stands for _________________ _________________ _________________ _________________

_________________.

Because valence electrons __________ each other, a molecule assumes a shape that keeps each valence electron pair as

_________________ _________________ from each other as possible.

Hence, VSEPR Theory predicts the _________________ of molecules.

In order to determine the shape of a molecule, you need to know:

1. How many atoms are bonded to the central atom.

2. The number of lone pairs attached to the central atom.

Lone pair = A _________________ pair of electrons.

Steric Number = The number of __________________ bonded to the central atom, plus the number of __________________

__________________.

To Determine the Shape of a Molecule

1. Draw Lewis Dot Structure 2. Determine steric number and number of lone pairs of central atom. 3. Use VSEPR Table to identify shape.

Polarity of Molecules

The VSEPR shape can be used to determine the polarity of a molecule.

A molecule with a __________________ three dimensional shape will be non-polar.

Testing for Polarity

Substances can be tested for their polarity by mixing them with liquids of known polarity. Generally, “___________________

__________________ ____________________” (two polar substances will be ____________________ and two non-polar

substances will be ____________________, but a polar and a non-polar substance will be ____________________)

Miscible: substances ___________________ each other (ex. salt and water) Immiscible: substances do ____________ dissolve each

other (ex. oil and water)

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Steric Number General Shape # Lone Pairs Specific Shape Picture Polar or Nonpolar?

1 Linear

2 Linear

Nonpolar

3 Trigonal Planar 0 Trigonal Planar

Nonpolar

3 Trigonal Planar 1 Bent

Polar

4 Tetrahedral 0 Tetrahedral

Nonpolar

4 Tetrahedral 1 Trigonal Pyramidal

Polar

4 Tetrahedral 2 Bent

Polar

5 Trigonal Bipyramidal 0 Trigonal Bipyramidal

Nonpolar

5 Trigonal Bipyramidal 1 Seesaw

Polar

5 Trigonal Bipyramidal 2 T-shaped

Polar

5 Trigonal Bipyramidal 3 Linear

Nonpolar

6 Octahedral 0 Octahedral

Nonpolar

6 Octahedral 1 Square Pyramidal

Polar

6 Octahedral 2 Square Planar

Nonpolar

29

Example: Determine the VSEPR shape and polarity of SF2.

Example Molecule

Lewis Structure Steric Number # lone pairs on central atom

Shape VSEPR Diagram Polar or Nonpolar?

CO2

BH3

CH4

NH3

H2O

PCl5

SF6

30

WS #7 (Learning Targets 3-15, 3-16, 3-17)

Complete the following table.

Example Molecule

Lewis Structure Steric Number

# lone pairs on central atom

VSEPR Diagram and Shape Polar or Nonpolar?

SiI4

SeF2

PBr5

BAt3

SBr6

NF3

CS2

SF6

31

Answers.

Molecule Lewis Structure Steric Number

# lone pairs on central atom

VSEPR Diagram and Shape Polar or Nonpolar?

SiI4

4 0

Nonpolar

SeF2

4 2

Polar

PBr5

5 0

Nonpolar

BAt3

3 0

Nonpolar

SBr6

6 0

Nonpolar

NF3

4 1

Polar

CS2

2 0

Nonpolar

32

WS #9 (Unit 3 Review)

1. (Learning Target 3-1) Discuss the development of the periodic table by Mendeleev.

What are the differences between Mendeleev’s Periodic Table and the Modern Periodic Table?

2. (Learning Target 3-2) Locate and state important properties of main chemical families including the alkali metals, alkaline

earth metals, transition metals, halogens, noble gases, lanthanides, actinides, and hydrogen.

Identify the period and group that each of the following elements belongs to, the number of valence electrons its neutral

atom has, and a property that all elements in that group share.

Element Period Group Valence Electrons Property

Ca

F

O

Ar

Fr

Ag xxx

U xxx

3. (Learning Target 3-3) Define atomic radius and use atomic structure to explain the periodic trend in atomic radius as you go

down a group on the Periodic Table and across a period.

4. (Learning Target 3-4) Define ionization energy and use atomic structure to explain the periodic trend in ionization energy as

you go down a group on the Periodic Table and across a period. List the two groups of elements that are exceptions to the

general trend and use orbital notation to explain why.

5. (Learning Target 3-5) Define electronegativity and use atomic structure to explain the periodic trend in electronegativity as

you go down a group on the Periodic Table and across a period. Which group of elements has zero electronegativity?

6. (Learning Target 3-6) Define ionic radius and compare the size of an anion to a neutral atom of the same element and a

cation to a neutral atom of the same element. Which type of elements from anions? Which type of elements form cations?

7. (Learning Target 3-6) Which has the larger radius, the neutral atom or the ion? Explain.

a. O or O2- b. K or K+

33

WS #9 - p. 2 – (con’t) (Unit 3 Review)

8. (Learning Targets 3-3, 3-4, 3-5, 3-6)

Complete the following table on Periodic Trends.

Electron Configuration (a) Na or S? Na: S:

(b) F or I? F: I:




9. (Learning Target 3-7) Draw electron dot diagrams for atoms, showing the correct number of valence electrons.

Draw Lewis Dot Structures for the following atoms or ions.

Element or Ion Lewis Dot Structure Element or Ion Lewis Dot Structure

Na+

H

Br

N3-

Ga

C

He

34

WS #9 – p. 3 (con’t) (Unit 3 Review)

10. (Learning Targets 3-8, 3-9, 3-14) Complete the following table.

Chemical Formula Electronegativity Difference

Type of Bond: Ionic, Polar Covalent or Non-Polar Covalent

Distribution of Charge (if Polar Covalent)

Lewis Dot Structure

CsBr

NF3

CaO

SO3

I2

BrF3

O3

11. (Learning Targets 3-15, 3-15) VSEPR Shapes

35

WS #9 – p.4 (con’t) (Unit 3 Review)

Complete the following table.

Molecule Lewis Dot Structure VSEPR Shape VSEPR Drawing Polar or Non-Polar?

AsCl3

SiH4

OBr2

GeO2

PAt5

SI6

BI3

12. (Learning Targets 3-11, 3-12, 3-13) Polyatomic Ions.

Determine the total number of electrons in each of the following polyatomic ions. Draw Lewis structures and indicate the formal

charge on each atom. Draw all possible resonance structures for the polyatomic ion.

(a) NO2- (b) BO3

- (c) C2O42-

13. (Learning Target 3-17) Define miscible and immiscible.

14. Identify two ways to tell whether atoms will form ionic or covalent bonds. Identify four properties of ionic compounds and

two properties of covalent compounds.

36

Answers.

(1) Mendeleev’s Periodic Table arranged elements by atomic mass, whereas the modern Periodic Table organizes elements by

atomic number. Mendeleev left gaps for elements whose properties were predicted based upon their location on his

Periodic Table. Mendeleev did not predict the existence of the noble gases which were discovered later.

(2)

Element Period Group Valence Electrons Property

Ca 4 Alkaline Earth Metals 2 Reactive metals; not found freely in nature

F 2 Halogens 7 “Salt-formers”; exist diatomically, exist as either solid, liquid or gas at room temperature

O 2 VIA 6 Non-metals

Ar 3 Noble Gases 8 Do not react to form compounds.

Fr 7 Alkali Metals 1 Highly reactive metals; not found freely in nature

Ag 5 Transition Metals 1 Metals: malleable, ductile and conductive

U 7 Inner Transition Metals/ Actinides

Varies Metals: malleable, ductile and conductive; Radioactive

(3) Atomic radius is the distance from the nucleus of an atom to the outermost electron. Atomic radius increases as you go

down a group, because, with each additional period, another energy level of electrons is being added. Atomic radius

decreases as you go across a group, because the number of protons in the nucleus increases, while no new energy levels

are being added, drawing the outermost electrons closer to the nucleus.

(4) Ionization energy is the energy required to remove the outermost electron. Ionization energy decreases as you go down a

group because, with each additional period, another energy level of electrons is being added, increasing the distance

between the outermost electrons and the nucleus, decreasing the attractive force. As you go across a period, the increase

in nuclear charge without the addition of more energy levels, increases the attractive force between the nucleus and its

outermost electrons, making the electrons harder to remove. The two groups that are exceptions to this trend are the

Boron group and the Oxygen group. With the Boron group, the removal of its outermost p1 electron leaves behind a stable

ion with a filled s-orbital, making that electron easier to remove than the outermost s2 electron in the alkaline earth metal

group. With the Oxygen group, the removal of its outermost p4 electron leaves behind a semi-stable half-filled p-orbital,

making that electron easier to remove than the outermost p3 electron in the nitrogen group.

(5) Electronegativity is the ability of a nucleus in a chemical bond to attract the electrons of the other atom toward itself.

Electronegativity decreases as you go down a group because, with each additional period, another energy level of electrons

is being added, increasing the distance between the outermost electrons and the nucleus, decreasing the attractive force.

As you go across a period, the increase in nuclear charge without the addition of more energy levels increases the attractive

force between the nucleus and its outermost electrons, making the nucleus better able to attract electrons toward itself.

The exception to this are the noble gases which have an electronegativity of zero because they do not form bonds.

(6) Metals form cations by losing electrons. Since electrons are lost, the cation is smaller in radius than its neutral atom.

Nonmetals form anions by gaining electrons. Since electrons are gained, the anion is larger in radius than its neutral atom.

(7) (a) O2- is larger than O because it has 2 more electrons. (b) K is larger than K+ because K has 1 more electron.

(8)

Electron Configuration (a) Na or S? Na: 1s2 2s2 2p6 3s1 S: 1s2 2s2 2p6 3s2 3p4

(b) F or I? F: 1s2 2s2 2p5 I: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p5


Na has a larger atomic radius because it has fewer protons than S and its outermost electrons are in the same energy level as sulfur’s. The fewer protons means that the Na

I has a larger atomic radius than F because its outermost electrons are in the 5th energy level, whereas F’s outermost electrons are only in the 2nd energy level.

37

has a smaller nuclear charge and exerts a smaller attractive force on the outermost electrons than S.


S has a higher ionization energy because it has a greater nuclear charge than Na, while its outermost electrons are in the same energy level as sulfur’s. This greater nuclear charge exerts a greater attractive force, requiring more energy to remove an electron.

F has a higher ionization energy than I. Since fluorine’s outermost electrons are closer to the nucleus than iodine’s, they experience a greater attractive force which requires more energy to remove them.


S has a higher electronegativity because it has a greater nuclear charge than Na, while its outermost electrons are in the same energy level as sulfur’s. This greater nuclear charge exerts a greater attractive force on its own and other atoms’ electrons in a bond.

F has a higher electronegativity than I. Since fluorine’s outermost electrons are closer to the nucleus than iodine’s, the nucleus is able to exert a greater attractive force on its own and other atom’s electrons in a bond.

(9)

Element or Ion

Lewis Dot Structure Element or Ion

Lewis Dot Structure

Na+ Na+ H

Br

N3-

Ga

C

He

(10)

Chemical Formula

Electronegativity Difference

Type of Bond: Ionic, Polar Covalent or Non-Polar Covalent

Distribution of Charge (if Polar Covalent)

Lewis Dot Structure

CsBr 2.1 Ionic N/A

NF3 1.0 Polar Covalent

38

CaO 2.5

Ionic N/A

SO3 1.0

Polar Covalent

I2 0 Non-Polar Covalent

N/A

BrF3 1.2

Polar Covalent

O3 0 Non-Polar

Covalent N/A

11)

Molecule Lewis Dot Structure VSEPR Shape VSEPR Drawing Polar or Non-Polar?

AsCl3

Trigonal Pyramidadl

Polar

SiH4

Tetrahedral

Non-Polar

OBr2

Bent

Polar

GeO2

Linear

Non-Polar

39

PAt5

Trigonal Bipyramidal

Non-Polar

SI6

Octahedral

Non-Polar

BI3

Trigonal Planar

Non-Polar

(12) (a) (b) (c)

(13) Miscible = completely soluble in each other; Immiscible = insoluble in each other

(14) Ionic compounds are between metals and nonmetals and ΔEN > 1.7. Covalent compounds are between non-metals only and

ΔEN <= 1.7. Ionic compounds form crystalline structures, are brittle, have high melting and boiling points and conduct electricity in

when molten or in solution. Covalent compounds have lower melting and boiling points and do not conduct electricity

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The Periodic Table Learning Targets (Your exam at the end of Unit 2 ...3+Packet+Honor… · Chemistry HP Unit 3 – The Periodic Table Learning Targets (Your exam at the end of Unit