Chemistry

Homework
 p. 193 #3-11, p. 207 - 208 # 31-32, 35-40, 55-56, 58-59,
62-63, 65 – DUE WEDNESDAY
 p. 199 # 16, 18-20, p. 207-208 # 42,45 – DUE
THURSDAY
 Quiz over Table 7.2 on Wends, Nov 9
 Test over Ch. 7 on Tuesday, Nov. 15 (tentative)


Review
1. What is an ion?
 ion – an atom with a positive or negative charge (an electron has been lost or gained)
2. Where on the periodic table are the metals located?
Nonmetals ?


Valence Electrons
Core Electrons – “inner” electrons, do not participate in bonding
 Valence Electrons – “outer” electrons, electrons in the highest occupied energy level
 number of valence electrons largely determines the chemical properties of an element
 So which electrons in an electron configuration are core and valence?
Be – [He] 2s 2 P – [Ne] 3s 2 3p 3
Core Valence Core Valence


The valence electrons an atom has is the same as its group number for a representative (main group) element


So how many valence electrons would Na have?
 1
 Iodine?
 7
 Krypton?
 8

 Valence electrons are VERY important for chemistry and reactions!
 Why?
 Chemists have electron dot structures – diagrams that show valence electrons as dots

 What do you notice about valence electrons down a group?
 Electron configuration remains the same down a group, only the core electrons change

Practice
 How many valence electrons does Al have? Draw the
Lewis dot structure.
 Electron configuration: [Ne] 3s 2 3s 1


Valence Electrons: 3
Lewis dot structure:

Crystal Structures
 In crystalline materials, elements repeat in repetition to give a 3D arrangement known as a crystal lattice
 A unit cell is the smallest part of the lattice that represents the entire lattice (repeated part)
 4 types of unit cells:
 Primitive
 Body-centered
 Face-centered
 Hexagonal close-packed

Crystal Structures
Primitive body-centered
Hexagonal close-packed face-centered

Part I: Metals
 Crystal coordination number (CN): number of neighbors to an atom in a unit cell
 Example: simple cubic is 6

Part I: Metals
 Crystal coordination number (CN): number of neighbors to an atom in a unit cell
 Face-centered:
 12
 Body-Centered:
 8
 Hexagonal:
 12

Let’s Think!
 We know noble gases are very stable. How can we get the other elements to become stable like the noble gases?
 Hint: Think of valence electrons!

Octet Rule
 All elements want eight valence electrons (except H and He)
 Why?
 Achieve noble gas configuration, most STABLE
 2 electrons come from the s subshell and 6 from the p subshell
 Why can H and He NOT have eight valence electrons
(Hint: they are considered stable with TWO valence electrons)
 No p subshell!

Octet Rule
 Elements will tend to gain or lose enough electrons to fulfill their octet
 Will metallic elements gain or lose electrons?
 Lose , much easier to lose electrons than gain electrons
 Will nonmetallic elements gains or lose electrons?
 Gain , much easier to gain electrons (can also share electrons)


Octet Rule
Metallic elements lose electrons
 Why?
 In simple terms…easier to lose than to gain to form the octet
 Nonmetallic elements gain electrons
 Why?
 In simple terms, easier to gain a few electrons than lose them
 Now let’s consider periodic trends!


Nuclear attraction across:
Shielding effect across:
Increases across, nonmetals held tighter than metals
Remains constant, does not offset nuclear attraction
 Kinetic energy across: Remains constant
 This means: Metals hold onto their electrons less tightly so it is easier for metals to lose their electrons. Additionally, nonmetals hold onto their electrons tighter . This trend, and the fact that nonmetals have higher electron affinity, means nonmetals tend to gain electrons.

Octet Rule
 If metals lose electrons, they will form ions. Will these ions be negatively or positively charged?
 If nonmetals gain electrons, will their ions be negatively or positively charged?


Octet Rule
Ions get special names to distinguish between those ions that are positively charged and those that are negatively charged
 Cation – positively charged ion
 Let’s try sodium!


Electron configuration:
[Ne] 3s 1
Original electron dot diagram:



How many electrons will it lose?
New electron dot diagram: Na +
New electron configuration
[Ne]
1


Octet Rule
Ions get special names to distinguish between those ions that are positively charged and those that are negatively charged
 Cation – positively charged ion


Octet Rule - Cations
Ions get special names to distinguish between those ions that are positively charged and those that are negatively charged
 Cation – positively charged ion

I wonder…
 Mg has the electron configuration of [Ne] 3s 2 . How many electrons would it lose or gain stability?
 2
 Will it take more or less energy to remove the second electron than the first electron?
 Let’s use electron dot structures to show its loss of 2e -

Octet Rule - Cations
 We just did one example of formation of a cation from
Group 1 and Group 2 . What can we say about the general formation of cations from
Group 1 and the general fromation of cations from
Group 2?
 Group 1 – form +1 cation
 Group 2 – form + 2 cation

 It’s extremely hard for transition metals to lose enough electrons to form a noble gas configuration. Why?
 They would have to lose all the s and d electrons!
 The transition metals ending with ns 1 (n-1)d 10 can lose that 1s electron to achieve pseudo noble-gas electron configuration
 What is one element that does this?
 Elements with the configuration ns 1 (n-1)d 5 can also lose that 1s electron to achieve half shell stability
 What is one element that can do this?

Octet Rule - Anions
 Ions get special names to distinguish between those ions that are positively charged and those that are negatively charged
 Anion – negatively charged ion
 Let’s try chlorine!
 Electron configuration: [Ne]3s 2 3p 5




Electron diagram:
How many electrons will it gain? 1
New electron diagram: Cl -
New electron configuration
[Ne]3s 2 3p 6

Octet Rule - Anions
 Ions get special names to distinguish between those ions that are positively charged and those that are negatively charged
 Anion – negatively charged ion

Octet Rule - Anions
 Ions get special names to distinguish between those ions that are positively charged and those that are negatively charged
 Anion – negatively charged ion

I wonder…
 O has the electron configuration of [He] 2s 2 2p 4 . How many electrons would it gain to gain stability?
 2
 Will it take more or less energy to gain the second electron than the first electron?
 Let’s use electron dot structures to show its gain of 2e -

Octet Rule - Anions
 How many electrons need to be gained by each group in group 5,6, and 7 to be stable?

Octet Rule

Common Ions
 Everyone loves memorizing stuff in chemistry! Let’s memorize some more!
 (Trust me, it’ll make your life 10x easier if you just memorize them now and get it over with!)

Common Ions

Let’s Think!
 We all know that certain elements will bond together to form molecules, such as NaCl (table salt). Why is this?
 Hint: Think of valence electrons.

Ionic Bonding
 Ionic compound – composed of metal cations and nonmetal anions
 They are still electrically neutral
 Why?
 The charges of the cation and anion cancel out
 The negatively charged anion and the positively charged cation attract each other by electrostatic forces (force holding a cation and anion together due to their charge)

Ionic Bonding
 Ionic compound – composed of metal cations and nonmetal anions
 Let’s go back to the formation of NaCl…
 What’s the electron configuration of Na?
[Ne] 3s 1
 Of Cl?

[Ne] 3s 2 3p 5
Their valence electrons can “add” together to form an ionic compound…now both elements are happy…just like Justin and Selena!

Ionic Bonding
 Ionic compound – composed of metal cations and nonmetal anions

Ionic Bonding
 Ionic compound – composed of metal cations and nonmetal anions

Ionic Bonding
 Ionic compound – composed of metal cations and nonmetal anions

Ionic Bonding
 So Na + and Cl combine to form NaCl
 NaCl is the chemical formula or a way to show the kinds and numbers of atoms in the smallest unit of a substance
 So there is 1 Na + for every 1 Cl -
 What would Mg +2 and Cl -1 form?


There is 1 Mg +2 for every 2 Cl -1 so MgCl
2
Think of it this way: the charges have to become neutral

Ionic Bonding
 The charges must become neutral
 What about Sr +2 and F ?
 SrF
2

Ionic Bonding
 Another way to do it is balance out the charges
 Ex:



1 Mg has a +2 charge and 1 O has a -2 charge. The charges are already balanced!
So we have….MgO
 What about….
1 Be has a +2 charge and 1 Br has a -1 charge.


How many Br will be needed to balance out the charges?
2…so that gives us….
 BeBr
2