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September 24 - October 6, 2017

9/13/2017

Housekeeping: Need to see progress on IA topics and thinking this week. If you need to speak to me outside of class, I'm available M-W between 10:35 - 1:00.

Your bonding quiz (sections 4.1-4.3) is this Friday.

Agenda:
Intermolecular Forces
Metallic Bonding
Review

Content Review:
Links:    Intermolecular Forces

Student Missions:
Mission 1: What's Keeping Us Together Other Than Love??
Mission Objectives. You should be able to...
1. List and describe the types of intermolecular forces
2. Deduce the types of intermolecular forces present in substances based on their structure and chemical formulas.
3. Explain the relationship between physical properties of covalent compounds in terms of their structure and intermolecular forces.

Covalent bonds holds atoms together within molecules, but what forces hold molecules together? The answer depends on the polarity and size of the molecules involved, and so said intermolecular forces will vary. The strength of intermolecular forces determine the physical properties of a substance. Volatility, solubility, and conductivity can all be predicted and explained from knowledge of the nature of intermolecular forces.

Dispersion forces (aka London forces): These are weak forces of attraction that occur between opposite ends of two temporary dipoles, usually with nonpolar molecules. A dipole occurs when the density of an electron cloud at any one moment be greater in one region of a molecule or atom. These dipoles are called temporary or instantaneous because they do not last long. When temporary dipoles influence the electron behavior of a nearby atom or molecule, induced dipoles result.

Polarizability is the ease of distortion of the electron cloud. Remember, most molecules that are asymmetric are polar. This means that they have a clear negative and positive end.

Factors that affect dispersion (London) forces:
1. number of valence electrons: greater the number of electrons, the larger the distance between valence electrons and the nucleus.
2. volume of electron cloud: large electron clouds can be polarized more easily.
3. molecule shape: large areas of interaction allow the London forces to have a greater magnitude, which leads to higher boiling points.

Dipole-dipole forces: Polar molecules that have a permanent separation of charge within their bonds as a result of electronegative differences have dipole-dipole attractions. Water has a clear positive end and a clear negative end. This is called a permanent dipole. When it bonds with another water molecule, a dipole-dipole attraction results. The strength varies depending on the distance and orientation of the dipoles.

The umbrella term van der Waals includes both dipole-dipole and dispersion forces.

Hydrogen bonding:  When a molecule contains hydrogen covalently bonded to a very electronegative atom (fluorine, oxygen, or nitrogen), these molecules are attracted to each other by a particularly strong force called a hydrogen bond. These are specific instances of dipole-dipole attraction. The large electronegativity difference results in the electron pair being pulled away from hydrogen. It now exerts a strong attractive force on a lone pair of electrons in a nearby molecule.

Hydrogen bonds are the strongest intermolecular force. As a result, they cause the boiling points of substances to be higher than normal.

Mission 2: Metallic Hook-Ups
Mission Objectives. You should be able to...
1. Explain the electrical conductivity and malleability in metals.
2. Explain the trends in melting points of metals.

Metallic bonding is the electrostatic attraction between a lattice of cations and delocalized electrons. Remember that delocalized electrons are not tied to a nucleus and free to move throughout the structure.

The strength of a metallic bond depends on three things: (1) number of valence electrons that can be delocalized, (2) charge of the metal ion, and (3) ionic radius of the metal cation.

Metals are good conductors of electricity because of the delocalized electrons. Charged species have to be free to move in order to conduct electricity. Metals are malleable, which means they can be hammered into shapes without breaking. Malleability is the result of cations sliding past each other, which leads to a rearrangement of the shape of the solid. The metallic bonds within the lattice structure do not have a defined direction; therefore, they are non-directional.

Melting points of metals depend on the strength of the attractive forces that hold the cations within the sea of delocalized electrons. The melting point of calcium is higher than potassium for the following reasons:

*Calcium has two delocalized electrons per atom and potassium has one. Therefore the electrostatic attraction between the cations and the delocalized electrons are greater.

**Calcium forms a 2+ cation and potassium forms a 1+. Therefore the electrostatic attraction between the positive ions and delocalized electrons is greater in calcium.

***Calcium has a smaller ionic radius that potassium. This implies that the delocalized electrons will be more strongly attracted to the Ca2+ ion.

Alloys, which are metallic solutions, can have a number of improved properties compared to their parent metallic element: increased strength, resistance to corrosion, enhanced magnetic properties, and greater ductility (ability to be drawn into wire). Adding additional atoms to the lattice reduces the ability of the cations to slide past each other and makes the metal stronger.

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September 11 - 20, 2017

9/11/2017

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Housekeeping: Need to see progress on IA topics and thinking this week. If you need to speak to me, I'm available M-W between 10:35 - 1:00.

Agenda:
1. VSEPR Theory & Molecular Geometry.

Content Review:
Links: VSEPR Theory Molecular Geometry

Student Missions:
Mission 1: What Molecules Look Like.
Mission Objectives. You should be able to...
1. Explain VSEPR Theory.
2. Draw Lewis structures of covalent compounds using VSEPR Theory.
3. Use VSEPR Theory to predict molecular geometry of different compounds
4. Use VSEPR Theory to predict molecular polarity.

VSEPR stands for Valence Shell Electron Pair Repulsion. This theory helps us to predict the shape of molecules using their valence electrons. There are six steps to drawing Lewis structures using the VSEPR theory. Mr. Anderson walks you through the process. We will work some practice problems in class.

Mission 2: Keeping Things Straight.
Mission Objectives. You should be able to...
1. Identify and sketch the shapes and bond angles of different covalent compounds.

Remember that the book (and therefore the exams) will refer to shared pairs of electrons as electron domains. Here is what we know so far:

Molecules (or species) with two electron domains tend to be linear with bond angles of 180 degrees.

Species with three electron domains tend to be trigonal planar with bond angles of 120 degrees. But if one of the electron domains is a lone pair (or unshared pair [two dots]), it doesn't show in the overall shape. This results in species that are bent and have bond angles of 117 degrees.

Species with four electron domains tend to be tetrahedral and have bond angles of 109.5 degrees.

Pages 108-109 provide excellent summary tables of these shapes.

Mission 3: Resonance.
Mission Objectives. You should be able to...
1. Draw and describe resonance structures of different compounds, such as benzene, carbonate, and ozone.
2. Explain the properties of giant covalent compounds in terms of their structures.
3. Define and explain coordinate covalent bonding.

Delocalization is when bonding electrons are not restricted to specific positions in molecules. They can spread out and give greater stability to thee molecule or ion. This happens frequently in molecules when there is more than one possible position for a double bond within a molecule. Example: Ozone. See below.

As a result, molecules like these demonstrate resonance. The Lewis structures are called resonance structures, as there are multiple versions of where the delocalized electrons can be. Look at the benzene example below. This is a specific example you are expected to know. You can look up carbonate on your own.

Resonance gives special properties to the structures where it occurs. It affects bond lengths and strengths, which influence reactivity. Acid and base strengths can be explained by resonance.

Play with this simulation.

Carbon has four allotropes: graphite, diamond, graphene and buckminsterfullerene. An allotrope is a structural modification of the same element. Each of the four examples contain nothing but carbon bonded in different ways. Graphite, diamond, and graphene are covalent network solids, in which the atoms are held together in a giant 3D lattice structure. Quartz is another example (SiO2). By contrast, buckminsterfullerene is molecular.

Read up on graphite, diamond, graphene, buckminsterfullerene and quartz. This is covered in pages 117-120. You're expected to know the properties of covalent network solids and details about each allotrope of carbon and silicon dioxide.

Coordinate covalent bonding. Typically, the shared pair of electrons originate from both atoms that form the bond; one atom contributes one electron to the shared pair and the second atom contributes the second electron. In coordinate covalent bonding, the shared pair of electrons comes from only one of the two atoms. Species that have coordinate covalent bonding: [NH4]+, [H3O]+, CO, Al2Cl6, and transition metal complexes. Look at the examples on pages 121-22.

Homework: Work on the related questions in your practice problem packet.

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September 4 - 15, 2017

9/3/2017

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Housekeeping: By Friday of this week, I expect to see three potential IA topics. No excuses. Check the "IA Topics" tab at the top right and click on the OWL link to see an exhaustive list of chemistry IA topics. Pick three and break each one down by your ability to figure out what kind of RQ you should ask, a potential hypothesis, and what kind of numeric data you hope to collect.

Agenda:
1. Discussion of ion formation.
2. How ionic/covalent bonds are formed.
3. How ionic/covalent compounds are named.
4. Writing ionic formulas.

Content Review:
Links: Chemical Bonding This page contains PowerPoints that cover ionic and covalent bonding and a summary table. This is a PowerPoint on Chemical Names & Formulas
Textbook Readings: Chapter 4.

Student Missions:
Mission 1: Ions, Ions, Ions!
Mission Objectives. By the end of this lesson, you should be able to:
1. Determine which elements form cations and which elements form anions.
2. Contemplate the structure of an ionic bond based on valence electrons.

Recall that the group number on the periodic table represents the number of valence electrons in that group. Also recall that group 18, the noble gases, have stable octets. The objective for any element that is not a noble gas is to lose or gain electrons to have a noble gas configuration. Stability is key.

Elements in G1 will LOSE their valence electrons to form ions with a +1 charge. Elements in G2 will LOSE theirs to form ions with a +2 charge. Elements in G13 will lose theirs to form ions with a +3 charge. Any time an element LOSES a valence electron, a cation is formed. Cations are positively charged ions. Metallic elements form cations.

Conversely, elements in G15 will GAIN 3 electrons to form ions with a -3 charge. Elements in G16 will GAIN 2 electrons to form ions with a -2 charge, and elements in G17 will GAIN 1 electron to form an ion with a -1 charge. Any time an element GAINS a valence electron, an anion is formed. Anions are negatively charged ions. Nonmetals form anions.

Transition metals have electron configurations that allow them to lose different numbers of electrons from their d subshell and form ions with different oxidation states. An oxidation state is the charge on an ion. The ions have different properties, such as forming compounds with different colors. An example would be the copper I and copper II ions ( Cu+ and Cu+2). One is green, the other blue.

Polyatomic ions are ions that are made up of several atoms covalently bonded, but have a positive or negative charge. I have given you a ion reference chart to use in class, but you need to memorize the different polyatomic ions, as they are not in the data booklet and you cannot use the ion chart on the test.

The images above (courtesy of wikipedia.org and http://sphweb.bumc.bu.edu) show the crystal lattice structure of salt (NaCl). The electrostatic attraction between cations and anions causes the ions to surround themselves with the opposite charged ion. The coordinate number is the number of surrounding ions. For salt, the coordination number is 6, as each Na+ is surrounded by 6 Cl-. Because these structures can get big, their formulas are merely the ratios of ions present. The simplest ratio is called a formula unit (fmu).

Lattice energy is a measure of the strength of attraction between ions within the lattice. The energy is greater for small, highly charged ions. Because of their lattice structures, ionic compounds have high melting and boiling points. The electrostatic attraction between the ions are strong and requires a lot of energy to break. Therefore, they are solids at room temperature and melt at very high temperatures.. When the charge is greater, the MPs and BPs are higher.

Volatility is the tendency of a substance to vaporize. Ionic compounds have a low volatility, or is non-existent. This means that they do not tend to vaporize because the bonds are too strong.

Solubility is the tendency of a substance to dissolve. Partial charges in water (a polar molecule) are attracted to opposite charges in an ionic compound and can dislodge the ions. When ions are surrounded by water, they become hydrated. Then the compound can dissolve. The state symbol for a compound dissolved in water is (aq) for aqueous. Nonpolar solvents do not have a charge, and therefore the ions do not pull apart. As a result, the substance is insoluble.

Conductivity & Brittleness. Based on an ability of a compound to conduct electricity, the ions have to be able to move. Therefore they cannot conduct electricity in the solid (s) state. The compound has to be molten (or liquid (l)) or dissolved in water (aq) to conduct electricity.

The ionic crystal lattice can shatter when a force is applied, so therefore it is brittle.

Homework: Re-read Mission 1. We will be working problems in class.

Mission 2: Sharing is Caring.
Mission Objectives. You should be able to...
1. Determine when covalent bonds are formed and which elements form them.
2. Predict the name of an ionic compound and a covalent compound.
3. Explain the difference between polar and nonpolar molecules.

Covalent bonds are the result of valence electrons being shared in order to achieve stable octets. If you think back...way back...to the short lesson on Lewis Dot, you'll recall that certain elements contained a certain number of dots that represented their valence electrons.

Anyhoo, if you take a look at groups 15, 16 & 17 (the nonmetals), this is where covalent bonding takes place. These elements can bond with themselves to share a pair (or two, or three) of electrons to form stable octets. Take a look below. Fluorine will share a pair of electrons with another fluorine and a single covalent bond is formed (represented by a line).

Group 16 elements will share two pairs of electrons (double covalent bond) and group 15 elements will share three pairs of electrons (triple covalent bond). Again, the goal is stability. The Octet Rule rules all.

Here is an interactive on covalent bonding. Play with it.

There are seven elements that are found in nature covalently bonded to each other. They are called diatomic elements. A useful memory aid is HONI Bring Fried Clams. So when you're writing these elements in chemical equations, you need to make sure you write them correctly.

Coordinate covalent bonds are bonds where one atom contributes both bonding electrons. You demonstrate them in structural formulas using an arrow that shows where the electrons came from. Once formed, the coordinate covalent bond is just like any other bond. Carbon monoxide is a good example of this.

Bond Disassociation Energy. This is the energy required to break covalent bonds. The unit is usually given in kJ/mol, which is the energy required to break one mole of bonds. Look below. Simply put, the larger the BDE, the stronger the covalent bond. Single bonds are easier to break than double bonds, and double bonds are easier to break than triple bonds, and so on and so forth. This will become important when we get to thermochemistry.

Mission 3: Electron Tug-of-War.
Mission Objectives. You should be able to...
1. Explain the difference between polar and nonpolar molecules.

Because covalent bonding involves sharing of electrons, there can be differences in how the electrons are shared between the bonded atoms. The character (or behavior) of the molecules depends on the kind of and number of atoms joined together, which in turn, determine molecular properties.

Below are pictures that demonstrate how the electrons are shared between atoms. If you think of it in terms of tug-of-war, there will either be an even pull on the electrons, or an uneven pull on the electrons. An even pull means that the molecule does not orient in any particular direction, which makes it symmetric and considered to be nonpolar. If there is an uneven pull on the electrons, then the molecule will orient in a particular direction (because the unshared pairs of electrons take up more space than shared pairs) and become asymmetric. These molecules are considered to be polar.

Homework: Complete the practice problems in the packet I gave you.

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G11 IB Chemistry Mad!Lab!

September 24 - October 6, 2017

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