focus learning targets for periodic trends and bonding...

Focus Learning Targets for Periodic Trends and Bonding

(1) Discuss the development of the periodic table by Mendeleev.

(2) Locate and state important properties of main chemical families including the alkali metals, alkaline

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

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

(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.

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

(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.

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

(8) Draw Lewis structures from chemical formulas.

(10) Calculate the total number of valence electrons in a polyatomic ion.

(11) Draw Lewis structures for polyatomic ions.

(12) Assign formal charges to atoms in polyatomic ions.

(13) Draw resonance structures for polyatomic ions.

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

(15) Assign shapes to molecules using VSEPR Theory and draw the VSEPR diagrams for a molecule

Periodic Table/ Periodic Trends ` Honors Chemistry

1. Indicate the different parts of the periodic table as instructed below: WS#1

a. Alkali metals blue

b. Alkaline earth metals green.

c. Transition metals orange.

d. representative metals

d. Noble gases red.

e..Halogens pink

f. Lanthanides purple.

g. Actinides yellow

2. Draw the Lewis dot structure for the elements 1-18 in the periodic table above.

3. What are valence electrons and why are they so important for understanding chemistry?

4. Identify each element as a metal, metalloid, or nonmetal.

a. chlorine

b. germanium

c. Tin

d. phosphorous

e. Nickel

f. Xenon

5. Give two examples of elements for each category.

a) noble gases

b) halogens

c) alkali metals

d) alkaline earth metals

6. Define atomic radius.

7. Circle the atom in each pair that has the largest atomic radius.

a) Al or C b) S or O

c) Mg or Na d) Sr or Ca

8. What trend in atomic radius do you see as you go down a group/family on the periodic table? What causes this trend?

9. What trend in atomic radius do you see as you go across a period/row on the periodic table? What causes this trend?

10. Define ionization energy? Circle the atom in each pair that has the greater ionization energy.

a) Li or B

b) Rb or K

c) Cl or S

d) Ca or Ba

e) P or S

f) Li or K

11. Define electronegativity.

12. What trend in electronegativity do you see as you go down a group/family on the periodic table? What causes this

trend?

13. What trend in electronegativity do you see as you go across a period/row on the periodic table? What causes this

trend?

14. Circle the atom in each pair that has the greater electronegativity.

a. K or Ga

b. Li or O

c. P or Cl

d. Br or As

e. Ca or Ba

f. O or S

15. Which species has the largest radius? Justify.

a. Mg or Mg2+ b. O or O2-

Chemical Bonds WS#2

1. Fill in the chart

Element Lewis Dot # of Valance e- Gain/Lose ___ e- Valance Charge

Na 1 Lose 1 +1

Be

Cl

S

Al

Ne

K

N

O

Ca

P

B

Mg

Lewis Dot Structures for Ionic compounds

Notes:

1. An ionic bond is an attraction of a cation for an anion resulting from the transfer of electrons. Remember, the smaller nonmetals are more electronegative and pull the electrons close, away from the larger, less electronegative metals.

2. When naming ionic compounds, the Metal is named first, followed by the nonmetal with an –ide ending. Ex. Sodium Fluorine becomes Sodium Fluoride.

3. Formula Unit: Lowest whole number ratio of elements in the compound. Ex. Ca3N2

1. Draw the Lewis Structure for Mg & Cl

Formula Unit: _________

Name of Compound:

2. Draw the Lewis Structure for Ca & S


Name of Compound:

3. Draw the Lewis Structure for Na & F


Name of Compound:

4. Draw the Lewis Structure for K & O


Name of Compound:

5. Draw the Lewis Structure for Be & N


Name of Compound:

6. Draw the Lewis Structure for Ca & P


Name of Compound:

7. Draw the Lewis Structure for Al & F


Name of Compound:

8. Draw the Lewis Structure for Ca & I


Name of Compound:

9. Draw the Lewis Structure for Rb & O


Name of Compound:

10. Draw the Lewis Structure for Sr & F


Name of Compound:

11. Draw the Lewis Structure for Al & Cl


Name of Compound:

12. Draw the Lewis Structure for Mg & P


Name of Compound:

13. Draw the Lewis Structure for B & O


Name of Compound:

14. Draw the Lewis Structure for Be & S


Name of Compound:

Lewis Dot Structures for Covalent Compounds WS#3

A chemical bond is an intramolecular (within the molecule) force holding two or more atoms together. Covalent

chemical bonds are formed by valence electrons being shared between two different atoms. Both atoms attain the noble

gas configuration with eight electrons (octet rule) or two electrons in their outer shell.

Octet Rule -Atoms bond in such a way that each atom acquires eight electrons in its outer shell.

Duet Rule - Hydrogen only requires 2 electrons to fill its outer shell and have He’s electron configuration. (1s2)

Lewis Dot Formula (also called an electron dot formula) – Shows the valence electrons, indicating the bonding between

atoms. The following guidelines will help draw the electron dot formulas correctly.

Rules for Writing Lewis Dot Structures

a. Count Valence electrons

b. Place least electronegative element in center

c. Give each atom 8 electrons except for hydrogen (only give hydrogen 2 electrons). Hydrogen cannot

go between 2 atoms, only around outside.

d. Compare number of dots with total valence electrons. Same number of dots and valence electrons

indicates all single bonds. Two dots too many indicates one double bond. Four dots too many

indicates 2 double bonds or one triple bond.

Exceptions to the Octet Rule

Boron tends to form compounds in which the boron atom has fewer than eight electrons around it.

Ex: BF3 – does not form a double bond because fluorine is much more electronegative than boron.

Some atoms exceed the octet rule. This is seen only for those elements in Period 3 and beyond. Ex:

SF6 ( Check class notes)

General comments about the Octet Rule

C,N,O, and F should always be assumed to obey the octet rule.

B and Be often have fewer than eight electrons around them in their compounds. These compounds

are very reactive.

The second-row elements never exceed the octet rule because they only have 2s and 2p orbitals.

Third-row and heavier elements often satisfy the octet rule but CAN exceed the rule by using their

empty valence d orbitals. (SF6)

When writing Lewis structures, satisfy the octet rule for the atoms first. If electrons remain after

this, then place them on the elements having available d orbitals (elements in period 3 or beyond)

1. Draw the electron dot structure for the following atoms:

a. Hydrogen b. Helium c. Silicon

d. Nitrogen e. Carbon f. Oxygen

g. Iodine h. Xenon i. Phosphorus

2. Draw the Lewis Structure for the following molecules

Lewis Structures: Single Bonds

a. H2

b. PH3

c. H2O

d. PCl3

e. NH3

f. CCl4

g. CH4

h. H2S

i. I2

j. CF4

3. Draw the Lewis Structure for the following molecules

Lewis Structures: Multiple Bonds

a. O2 b. C2H4

c. CO2

d. HCN

e. N2

f. C2H2

g. CH2O

h. CHCl3

i. XeF2

j. SF4

4. Draw the Lewis Structure for the following polyatomic ions

a. CN-1

b. H3O+1

c. NH4+

d. NO3-

e. CO32-

f. SO42-

g. ClO3-

h. NO2-

Formal Charges: WS#3

Formal charge can be a useful concept in determining the most likely Lewis structure for a compound. Sometimes

when you have more than one reasonable Lewis dot structure formal charge can help you to determine the most

stable Lewis dot structure. The Lewis dot structure that minimizes the charges on each atom is the correct one.

The most stable structure has:

• the lowest possible formal charge on each atom,

• the most negative formal charge is on the most electronegative atoms.

Formal charge on an atom= number of valence electrons in free atom - number of valence electron in

bonded atom

1. Identify the formal charges on each element in these structures

2.

3.

Resonance: Sometimes more than one valid Lewis structure is possible for a given molecule. Resonance

shows that electrons are not localized to one atom but instead travel throughout the molecule. Resonance

structures got their name because scientists originally thought that the bonds would switch or resonate between

the different positions. Further research revealed that compounds that can be written with resonance structures

actually have bonds that are hybrids of the other bonds in the compound.

4. Draw the Lewis dot structure for the following compounds and draw all three possible resonance structure for

the following compound. Assign formal charges in each and decide which one is the best structure.

a. PO43-

b. COCl2

c. NO3-

d. SO42-

e. CO32-

VSEPR (Valence Shell Electron Pair Repulsion) WS#4

The valence shell electron pair repulsion (VSEPR) theory was developed as a way to predict molecular geometries

based on Lewis electron dot diagrams. The molecular geometry of a molecule influences its physical properties,

chemical properties, and biological properties. Molecular geometry is associated with the chemistry of vision, smell

and odors, taste, drug reactions and enzyme controlled reactions. As you learned in Honors Biology, most enzymes

will react with molecules possessing only a certain, specific shape.

VSEPR theory states that because electron pairs repel, molecules adjust their shapes so that the valence-electron

pairs are as far apart as possible.

Fill in the table using the VSEPR chart provided in the class.

Molecule Total#

valence

electrons

Lewis Dot Structure Total #

electron

pairs

Elctron

geometry

Bonding

pair

Lone

Pairs

Shape of

the

molecule

ICl

BrO3-

SO3

CH2O

SF4

PF5

Polarity & Electronegativity

When two atoms combine, the difference between their electro negativities is an indication of the type of bond that will

form. If the difference between the electronegativities of the two atoms is small, neither atom can take the shared electrons

completely away from the other atom, and the resulting bond will be covalent. If the difference between the electro

negativities is large, the more electronegative atom will take the bonding electrons completely away from the other atom

(electron transfer will occur), and the bond will be ionic.

Classifying Bonds Using Electronegativity Differences

Electronegativity Difference

Bond Type

0 - 0.2

Nonpolar covalent bond

0.3 - 1.7

Polar covalent bond

1.8

Ionic bond

1. What is the difference between a covalent, polar covalent and ionic bond?

2. Complete the table by predicting the type of bond that will form between the following elements only using the

periodic table.

Elements Bond Type- Ionic or Covalent?

Beryllium and Flourine

Selenium and Chlorine

Strontium and chlorine

nitrogen and Iodine

Sulfur and Phosphrous

Magnesium and oxygen

3. For each of the following molecules, determine if it is covalent, polar covalent, or ionic. Show your work by

listing the electronegativities of each element in the bond.

Molecule Electronegativity

Values

Difference in

Electronegativity

Bond Type? Ionic,

Polar Covalent,

Nonpolar Covalent?

H – Cl H:

Cl:

H – H H:

H:

Cl - Cl Cl:

Cl:

C – O C:

O:

Ca – O Ca:

O:

Al – F Al:

F:

³

focus learning targets for periodic trends and bonding...

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