Honors Chemistry 1: Chapter 4 Study · Web viewHonors Chemistry 1: Chapter 4 Study Guide 24 | Page DAY 1: Section 4.1 Before coming to class, you should have: Read section 4.1 Completed

Honors Chemistry 1: Chapter 4 Study Guide

DAY 1: Section 4.1

Before coming to class, you should have:

Read section 4.1 Completed warm-up practice problems 1-4 Paid special attention to trends on the periodic table

Key concepts:

Law of octaves

Periodic Law

Valance Electrons

Group

Period

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Warm-Up Practice Problems:

1) Write the electron configuration AND create an orbital diagram for each of the listed elementsPut your orbital diagram first, and then write the configuration to the right.

a. Lithium

b. Sodium

c. Potassium

d. Rubidium

2) Look at all of your electron configurations in question 1. What similarities do you see for those four elements?

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3) Write the electron configuration AND create an orbital diagram for each of the listed elements

a. Fluorine

b. Chlorine

c. Bromine

d. Iodine

4) Look at all of your electron configurations in question 3. What similarities do you see for those four elements?

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5) Write the electron configuration AND create an orbital diagram for each of the listed elementsa. Carbon

b. Silicon

c. Germanium

d. Tin

e. Lead

6) Look at all of your electron configurations in question 5. What similarities do you see for those five elements?

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Figure 4.1a: Mendeleev’s Predictions

How did Mendeleev organize his version of the periodic table?

How did the “Law of Octaves” help Mendeleev organize his period table?

Why were these “gaps” present in Mendeleev’s periodic table to begin with?

How did Mendeleev arrive at these predictions?

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Figure 4.1b: Blocks of the periodic table

Using the information in figure 4.1b, list the number of valance electrons for each of the following elements:

Sodium: _____ Sulfur: _____ Oxygen: _____ Fluorine: _____

Magnesium: _____ Aluminum: _____ Silicon: _____ Carbon: _____

Lithium: _____ Tin: _____ Boron: _____ Chlorine: _____

Iodine: _____ Calcium: _____ Neon: _____ Argon: _____

Barium: _____ Strontium: _____ Krypton: _____ Potassium: _____

Arsenic: _____ Antimony: _____ Lead: _____ Xenon: _____

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Predicting:

I’ve repeatedly said that all main group elements want to be noble gases, which is the most stable group of atoms on the periodic table (Think of this as “noble gas envy”). Predict the formula for the following ionic compounds, and think in terms of helping each of these atoms find a way to have the electron configuration of a noble gas:

1) Sodium bonding with Chlorine

2) Sodium bonding with Bromine

3) Sodium bonding with Iodine

4) Lithium bonding with fluorine

5) Sodium bonding with fluorine

6) Potassium bonding with fluorine

7) Calcium bonding with oxygen

8) Sodium bonding with oxygen

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Day 2: Section 4.2 – Touring The Periodic Table


Read section 4.2 Reviewed section 4.1

Key concepts:

Main-Group Elements

Alkali Metals

Alkaline Earth Metals

Halogens

Noble Gases

Transition Metals

Properties of metals

o Ductile

o Malleable

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Lanthanides

Actinides

Alloys

At the top of each group, label the ionic charge for each element within that group.

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Figure 4.2a: Looking for trends among the Alkali Metals

Answer the following:

1) How does the melting point change as the atomic number increases for alkali metals?

2) How does the density change as the atomic number increases for alkali metals?

3) How does the atomic radius change as the atomic number increases for alkali metals?

4) How does metallic hardness change as the atomic number increases for alkali metals?

5) Which element shown as the greatest temperature range in its liquid state?

6) What is the difference in melting points between sodium and potassium?

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Figure 4.2b: Know the differences between METALS and NONMETALS

List a variety of general properties for METALS and NONMETALS:

METALS:

NONMETALS:

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Day 3: Section 4.3—Periodic Trends


Read section 4.3 Reviewed sections 4.1 and 4.2

Key Concepts:

Periodic trend

Ionization Energy

Atomic Radius

Bond Length

Electronegativity

Electron affinity

Boiling/Melting Point

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Figure 4.3a: What is ionization energy?

Important fact about ionization:

IN GENERAL, METAL ATOMS TEND TO ___________________ ELECTRONS, AND NONMETAL ATOMS

TEND TO ______________ ELECTRONS!!!!

Cations and Anions:

Cation:

Anion:

Examples:

Write the ionic charge for each of the given atoms once it has ionized

Sodium: _____ Fluorine: _____ Oxygen: _____ Nitrogen: ____

Sulfur: _____ Magnesium: _____ Calcium: _____ Aluminum: _____

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List the protons and electrons for each ion listed below:

ION Protons Electrons ION Protons ElectronsCa2+ Cu+

P3- N3-

Br - Au+

K+ Ba2+

Ag + Cl -

S2- H+

Al3+ Na+

Fe3+ Cs+

F - B3+

Figure 4.3b: Trends in Ionization Energy

Define ionization energy (do it again, even if you have already…it’s very important!!):

Explain WHY the trend shown in figure 4.3b exists as you move DOWN groups:

Explain WHY the trend shown in figure 4.3b exists as you move ACROSS periods:

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Figure 4.3c: Ionization energy for main-block elements

Generally speaking, as the group number increases, the ionization energy ________________________

As the atomic number increases WITHIN A GROUP, the ionization energy _________________________

Why do the noble gases have such incredibly high ionization energies compared to the rest of the elements shown in this graph? How does this describe their reactivity?

Which group (give the name…e.g., halogens, noble gases, alkaline earth metals, alkali metals) has the lowest ionization energy? What does this mean in terms of their reactivity?

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Figure 4.3d: Trends in atomic radius

Explain the atomic radius trend as you go down a group:

Explain the atomic radius trend as you go across a period:

Figure 4.3e: Plotting group number as a function of atomic radius

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Figure 4.3f: Period trends in IONIC radius

Figure 4.3g: Bond Radius

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Why is chlorine’s bond radius shorter than iodine’s bond radius?

ELECTRONEGATIVITY:

Figure 4.3h: Electronegativity Trends

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Figure 4.3i: Periodic Trends in electron affinity

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Explain why the halogens, out of ALL OF THE CHEMICAL GROUPS, have the highest electron affinity:

Why is the electron affinity higher in alkali metals than in alkaline earth metals?

Figure 4.3j: Periodic Trends in Melting/Boiling Point

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Of all of the elements shown, which seems to be the “biggest exception” to the general trend in period 6?

Why might tungsten have the highest melting and boiling point of all elements in period 6?

Approximate the following:

Au boiling point: ______ Pb melting point: _____ Hg freezing point: _______

Temp range that barium is a liquid: __________________

Section 4.4—Where did the elements come from?

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Read section 4.4 Reviewed sections 4.1-4.3

Key Concepts:

Big Bang Theory

Supernova

Nuclear Fission

Nuclear Fusion

Figure 4.4a: How elements are fused inside of stars

SUMMARY OF PERIODIC TRENDS

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Order the following elements based on the trend listed:

Electronegativity: Pb, Cl, Ba, F, Ca

Atomic Radius: C, Li, F, N, O

Atomic Radius: Ca, Ba, Sr, Mg, Be

Electron Affinity: Cl, I, Br, F

Electronegativity: K, Li, Rb, Cs, Na

Ionization energy: Sb, N, Bi, P, As

Atomic & Ionic Radii for common cations and anions

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Honors Chemistry 1: Chapter 4 Study · Web viewHonors Chemistry 1: Chapter 4 Study Guide 24 | Page DAY 1: Section 4.1 Before coming to class, you should have: Read section 4.1 Completed