Bellwork 9-5-20141. Analyze the Bohr model on the piece of

butcher paper on your desk. 2. 5 min: Write down everything you know

about the element in your own section.• Valence electrons, reactivity, group, period etc.

3. 5 min: Talk with your group and come to a consensus on the information you want to include in the center square

4. Present center information to the class- Pick a spokesperson for this!

Bellwork: Honors 9-4-2014 1. With your group, take 5 minutes to finalize

plans for you video2. You will only have 10 minutes to film and

get it right! • Don’t forget to include everything on the note

taker (valence e- and elements for groups 1, 2, 7, 8) and one thing you researched!

3. Designate 1 person in your group to get a lap top and sign it out

4. BE CAREFUL WITH LAP TOPS WHEN FILMING!

Butcher Paper Activity1. Analyze the Bohr model on the piece of

butcher paper on your desk. 2. 5 min: Write down everything you know

about the element in your own section.• Valence electrons, reactivity, group, period etc.

3. 5 min: Talk with your group and come to a consensus on the information you want to include in the center square

4. Present center information to the class- Pick a spokesperson for this!

Objective• You will be able to analyze a Bohr model

of a specific element and determine all properties of that element that we have discussed (valence electrons, reactivity, group, periodic table trend)

Periodic Table Trends1. Atomic radius 2. Ionization energy 3. Electronegativity

- As you go up in period #, you add an energy shell of electrons or another orbital- As you go up in group #, you add an electron to the outer electron shell (excluding the transition metals)

Number on top of each group = valence electrons (ignore the 1 for groups 13-18)

On Your Butcher Paper•Write down how many energy shells (orbitals) the element has

- Keep in mind, these are shells that CONTAIN electrons

Real model of the atom

1. Atomic RadiusAtomic radius• The distance from the nucleus to the

outmost electrons depending on pull from protons

Thought question:• But the electrons can be anywhere as

we just discussed in the last slide… so how can we predict where the outermost electrons are?

Atomic RadiusExample 1 – Right Left: • Compare Potassium (K) and Krypton (Kr)• Which do you think has a bigger atomic

radius- Period (outermost shell)- Valence electrons- # of Protons

Talk with your neighbor

Atomic Radius

Atomic Radius

Potassium Krypton

Which has a larger atomic radius?

Atomic RadiusAnswer:Shells: They are in the same Period so they both have the same number of shells (orbitals) Valence e-: Kr has more valence shell electrons but they are in the same orbital as KProtons: Kr has more protons in the center than K• Protons pull on electrons so the electrons in

Kr are being pulled inward moreResult: Kr is smaller than K

Atomic Radius Conclusion Ex. 1 – Right Left

• Atomic radius increases as you go right left• Because the fewer protons there are,

the smaller the pull on electrons allowing radius to be larger

Mark on your table at the bottom

Atomic Radius Example 2: Top Bottom• Compare Lithium (Li) and Cesium (Cs)• Which has a bigger atomic radius and

why? Talk with a neighbor

Atomic Radius

Atomic Radius

Cesium Which has a larger atomic radius?

Atomic Radius Answer: • Shells: Li has fewer electron shells than

Cs• Valence e-: Li and Cs have the same

number of valence electrons • Protons: Cs has more protons so more

pull but it they cannot reach the outer shells as well

• Result: Li is smaller than Cs

Atomic Radius Conclusion: • Atomic radius increases as you go

DOWN the periodic table • Because added energy shells (orbitals)

increases the distance of electrons from the nucleus

• Mark on your table to the left

Practice with a neighbor 1. Which element has a larger atomic

radius, Strontium (Sr) or Silver (Ag)? WHY?

2. Which element has a SMALLER atomic radius, Magnesium (Mg) or Rubidium (Rb)? WHY?

Strontium – less protons, less pull on e- more making atom bigger

Magnesium – fewer shells, e- closer to nucleus

On Your Butcher Paper•Is the atomic radius of your element big or small? Choose an element to compare it to (using left/right or up/down trends)• Explain to your group.

Bellwork 9-9-20141. Get out homework packet on trends, I will

be stamping EACH section (3 stamps total)2. What group on the periodic table is

Francium (Fr) in?3. How many valence electrons does Francium

(Fr) have? 4. How many energy shells does Francium (Fr)

have?

**Pass back homework packets

Element of the Day – Hydrogen •http://www.periodicvideos.com

2. Ionization Energy Ionization Energy:• The amount of energy required to

remove one electron from the outermost shell of an element

Question:• What do you think is most important to

consider when determining ionization energy of an element?

2. Ionization Energy Example 1: Left Right• Which element has a higher ionization

energy, Lithium or Fluorine? Why? - Think about valence electrons

Talk with your neighbor

Ionization Energy

Ionization Energy Which has a higher ionization energy?

Ionization Energy Answer:• Valence electrons: Fluorine almost has

a full outer shell of electrons and does not want to give them up. Lithium has only one and will willingly give it up to Fluorine

• Result: Fluorine has a higher ionization than Lithium

Ionization Energy Conclusion:• Ionization energy increases as you

go left right- Because it gets harder to pull electrons

away from elements with almost full outer shells Mark on your table on the top

Ionization Energy Example 2: Bottom Top • Which has a higher ionization energy,

Iodine (I) or Fluorine (F)? Why? - Think about atomic radius

Talk with a neighbor

Ionization Energy

Ionization Energy

Ionization Energy Answer:• Valence electrons: same in both• Atomic radius: iodine has a larger

atomic radius than fluorine so electrons are farther away from the nucleus and thus have a weaker pull on them from the nucleus

• Result: Fluorine has a higher ionization energy

Ionization Energy Conclusion:• Ionization energy increases as you

go from bottom to top - Because it gets harder to pull away an

electron as they get closer to the nucleus due to the pull of protonsMark on your table on the right

Practice with a neighbor 1. Which element has a higher ionization

energy, Carbon (C) or Neon (Ne)? WHY?

2. Which element has a SMALLER ionization energy, Cesium (Cs) or Sodium (Na)? WHY?

Neon – full outer shell of e-, harder to remove e-

Cesium – more e- shells, outer electrons easier to remove

On Your Butcher Paper•Pick an element on the periodic table in the same period OR group as your element • Is the ionization of your element HIGH or LOW compared to the one you chose? Explain to your group

3. Electronegativity Electronegativity• The amount of attraction the element

has for electrons (how much it hogs e-)• Also called ELECTRON AFFINITY – how

badly an element wants electrons to eventually become stable

Thought Question • What do you think determines an

elements “desire” for electrons?

3. Electronegativity Example 1: Left Right• Which is more electronegative, sodium

(Na) or chlorine (Cl) and why? - Think about valence electrons

Talk with your neighbor

Electronegativity

Electronegativity

Sodium Chlorine

ElectronegativityAnswer:• Valence electrons: Chlorine has more

valence shell electrons (7 e-) than sodium (1 e-) and therefore wants more electrons to fill its shell and become stable

Conclusion:• Electronegativity increases as you go

from left right- Because elements have more electrons in their

valence shell and WANT more to become stable (full outer shell)Mark on your table on the top

ElectronegativityExample 2: Bottom Top • Is Selenium (Se) more or less

electronegative than Oxygen (O)? Why? - Think about atom radius and number

of valence shells Talk with a neighbor

Electronegativity

ElectronegativitySelenium

Electronegativity Answer:• Valence electrons: all have 1 valence electrons• Atomic radius: Barium has more electron shells

and therefore electrons are farther away from the nucleus and not held as tightly. Beryllium has fewer electron shells and therefore holds electrons more tightly so it WANTS them more.

Conclusion:• Electronegativity increases as you go from

bottom top - Because elements have a smaller radius at the top and hold onto electrons more tightly (want them more)Mark on your table on the right

Practice with a neighbor 1. Which element has a higher

electronegativity (electron affinity), Argon (Ar) or Magnesium (Mg)? WHY?

2. Which element has a LOWER electronegativity, Xenon (Xe) or Neon (Ne)?

Argon – full outer shell of e-, does not want to give up

Xenon – more electron shell, outer e- not held as tightly (doesn’t want them as much)

On Your Butcher Paper •Choose an element on the periodic table in the same period OR group as your element. • Is the electronegativity of your element HIGHER or LOWER compared to that element? Explain to your neighbors.

Topics to be covered on quiz Wed.1. Average atomic mass (honors only)2. History of periodic table – from notes • Who came up with the table we use today and how did he

construct it?

3. Periodic table groups – from note-taker• Given a Bohr model, tell me what group, reactivity, # of

valence e-

4. Periodic table trends – from notes and HW

• Given two elements, which is more electronegative? Why? • Given two elements, which has a larger atomic radius? Why?• Given two elements, which has a higher ionization energy?

Why?

Which element has a SMALLER atomic radius? WHY?

Br K

Which element has a LARGER atomic radius? WHY?

SPo

Which element is more electronegative? WHY?

C F

Which element has a LOWER electronegativity? WHY?

Xe He

Which element is has a HIGHER ionization energy? WHY?

P Mg

Which element has a LOWER ionization energy? WHY?

Cs K

Exit Slip 9-5-20141. What are two things you need to look at

when determining atomic radius?2. What are two things you need to look at

when determining ionization energy?3. What are two things you need to look at

when determining electronegativity?

Homework – COMPLETE trends packet due WednesdayQuiz Wednesday

# of protons, # of energy shells

# of valence electrons, atomic radius (pull of protons)

# of valence electrons, atomic radius (pull of protons)

Steps to Solving Trend Problems Atomic Radius – size of atom 1. Find the elements on the periodic table2. If Left – Right look at # of protonsa. More protons = more pull = smaller

radiusb. Less protons = less pull = larger radius

3. If Up – Down look at # of shellsa. More shells = larger radiusb. Less shells = smaller radius

Steps to Solving Trend Problems Ionization Energy – energy needed to remove e-1. Find the elements on the periodic table2. If Left – Right look at # of valence electronsa. More valence electrons = harder to pull away =

higher ionization energy b. Less valence electrons = easier to pull away =

lower ionization energy 3. If Up – Down look at # of shellsa. More shells = lower ionization energy because

electrons farther from nucleus (easier to remove) b. Less shells = higher ionization energy because

electrons closer to nucleus (harder to remove)

Steps to Solving Trend Problems Electronegativity – desire for or attraction to electrons 1. Find the elements on the periodic table2. If Left – Right look at # of valence electronsa. More valence electrons = more attracted to

electrons = higher electronegativityb. Less valence electrons = less attracted to electrons

= lower electronegativity3. If Up – Down look at # of shellsa. More shells = lower electronegativity because

electrons farther from nucleus (not as attracted)b. Less shells = higher electronegativity because

electrons closer to nucleus (more attracted)