Unit 9 Bonding

CH4 methane gas molecule The attractive forces between atoms leads to chemical bonds that result in chemical compounds. Ask the students how the valence electrons are shared in methane molecule. You may need to remind the students that H has one valence electron while C has 4 valence electrons.

Why do atoms form bonds? Atoms form bonds due to the need to have the most stable configuration for its electrons. Tell the students that in order for sodium chloride to form, sodium atom loses its valence electron while chlorine atom (being more electronegative) gains the electron from sodium. Atoms lose, gain, or share valence electrons in order to achieve a lower energy state (stable).

How atoms bond with each other depends on: Electronegativity Ionization Energy # Valence Electrons e I want an electron e e e

Metallic bonding Metallic bonds are formed by metal atoms. Ask the student what they remember about the electronegativity and ionization energies of metals. (Low electronegativity and low ionization energy); The valence electrons of metals are easily displaced or removed or delocalized. http://www.launc.tased.edu.au/online/sciences/PhysSci/pschem/metals/Metals.htm

Metallic bonding Metallic bonding is the strong attraction between closely packed positive metal ions and a 'sea' of delocalized electrons. What are the properties of metals? How does metallic bonding affect the properties of metals? The attraction between the metal ions and the delocalized electrons must be overcome to melt or to boil a metal. Some of the attractions must be overcome to melt a metal and all of them must be overcome to boil it. These attractive forces are strong, so metals have high melting and boiling points. The delocalized electrons are also able to move through the metal structure. When a potential difference is applied, the electrons will move together, allowing an electric current to flow through the metal. Delocalized electrons make metals malleable and ductile. http://www.bbc.co.uk/schools/gcsebitesize Watch This

Metallic bonds What property of metals is illustrated above? The illustration shows the malleability of metals. The atoms in metals slide past each other and held tightly by free-moving (sea of) electrons. What property of metals is illustrated above?

Ionic Bonding an electrostatic force Electrostatic refers to the attraction between opposite charges Stronger than metallic bonds because of the opposite charges

Ionic bond Formed by the transfer of electrons between a metal and a nonmetal # of e- lost by metal = # of e- gained by nonmetal

Charge The number of electrons that need to be lost or gained by an atom so it has the same electronic configuration as a noble gas. When an atom gains or loses electrons, it is called an ION. + ions are cations - ions are anions The charge of the ion is called the OXIDATION NUMBER.

Watch This HAVE STUDENTS LABEL OXICATION NUMBERS ON PT

Charge based on “copy cat” principle Sulfur wants to be like Argon (1s22s22p63s23p6) Sulfur Atom 1s22s22p63s23p4 16 Protons 16(+) 16 Electrons 16 (-) 0 No charge Sulfur has 6 valence e- and will gain 2 more to complete an octet. The result is a: Sulfur Anion 1s22s22p63s23p6 16 Protons 16 (+) 18 Electrons 18 (-) 2- Charge

Ionic Bonds Strong electrostatic (positive-negative) force in ionic compounds makes a strong ionic bond. How does the strong ionic bond affect the properties of ionic compounds? Tell the students that it requires a lot of energy (high temperature) to break an ionic solid (or ionic bonds) into its component ions. As a result ionic compounds have: a. High melting and boiling points b. Brittle; the substance breaks into smaller units when hammered or turns into a powder Ionic compounds also conduct electricity when dissolved in water (electrolytes) or when molten. It does not carry electrical charges (conduct electricity) in solid state. Formula unit – smallest unit of an ionic compound; lowest whole number ratio of ions represented in an ionic compound

The greater the difference between eN values of 2 atoms is, the more ionic the bond will be. We call this “ionic character”. 0.8- 4.0 = 3.2 very ionic

Ionic Bond Na+ Cl- Show the students the different representations for the formation of an ionic bond. Use the model to the left to show how metals lose electrons to nonmetals. Use the model to the right to illustrate how ionic bonds are represented as Lewis dot structures. http://www.chm.bris.ac.uk/pt/harvey/gcse/ionic.htm

Writing Lewis Dot Structures Element symbol represents the kernel (core) of the atom (nucleus and inner e-) Dots represent the valence e- www.meta-synthesis.com 

Writing Lewis Dot Structures - Ionic Metals tend to lose e- while nonmetals tend to gain electrons Illustrate the formation of NaCl from Na and Cl atoms. Point to the lost of an Na electron forming Na+ while Cl gained the electron forming Cl-. Have the students explain the formation of MgO and CaCl2 using the next two representaitons. Ionic bonds hyperphysics.phy-astr.gsu.edu

Writing Lewis Dot Structures – Ionic Bonds Metals tend to lose e- while nonmetals tend to gain electrons Illustrate how to write the Lewis Dot Structure of CaCl2. Emphasize the writing of brackets [] and ionic charges. chemistry58.wikispaces.com Ionic bonds

Properties of ionic compounds (all related to the strong attraction between the + and – charges) high melting points and boiling points hard solids good conductors – in aqueous solutions and when molten have a crystal lattice structure Ions are here

Really, we don’t hate you.

Covalent Bond Covalent Bond –formed when two nonmetals share pairs of valence electrons in order to obtain the electron configuration of a noble gas Molecule - formed when two or more atoms bond covalently. (A molecule is to a covalent bond as a formula unit is to an ionic bond.)

How covalent atoms bond

Diatomic Molecules HOFBrINCl Share electrons when they bond together

Polyatomic Ions covalently bonded group of atoms, with a charge Watch this They are listed on your STAAR chart. You will not have a bad time.

Properties of Covalent Molecules Can exist as gases, liquids, or solids depending on molecular mass and polarity Usually have lower MP and BP than ionic compounds of the same mass Do not usually dissociate (break apart into ions) in water Do not conduct electricity

How to draw Lewis dot structures for covalent molecules. Write the formula for the compound. Count the total number of valence electrons. Predict the location of the atoms: Hydrogen is NEVER the central atom. If carbon is present, it is ALWAYS the central atom. If there is only 1 atom of an element, it is the central atom. The least electronegative atom is generally the central atom. Place one electron PAIR between the central atom and each ligand (side atom) to “hook” the atoms together. Dot the remaining electrons in pairs around the compound to complete the octet. Start with the ligands. Check that each atom has an octet. (H only needs a pair, not an octet.) Watch This

Lewis Structures for Molecules Draw the Lewis dot structure for these molecules: Hydrogen + Bromine (HBr) Carbon + Chlorine (CCl4)

Writing Lewis Dot Structures - Covalent Bonds Bonding e- Pairs Lone Pairs (nonbonding electrons) Identify and describe to the class the lone pairs (non bonding electrons) and the bonding electron pairs or ligands. Provide more examples to your L-classes, if necessary. Covalent bonds

Number of bonds Single Bonds - when one pair of e- is shared between atoms Double bond – when atoms share 2 pairs of valence electrons; ex. O2 Triple bond – when atoms share 3 pairs of valence electrons; ex. N2

Describing bonds Sigma bond - the first bond between 2 atoms A single bond is a sigma bond. Pi bond - the second bond between 2 atoms A double bond consists of a sigma bond and a pi bond. A triple bond consists of a sigma bond and two pi bonds.

Carbon can form single, double and triple bonds with itself.

Why are molecular shapes important? The shape of a molecule plays a very important role in determining its properties. Molecular shapes determine the properties of a substance. For example, hemoglobin performs an important function in respiration (transport of O2 and CO2). Sickle cell anemia results from malformed hemoglobin resulting to a change in the molecular geometry of hemoglobin, restricting the normal function of hemoglobin. Properties such as smell, taste, and proper targeting (of drugs) are all the result of molecular shape.

VSEPR Theory also called electron geometry Electron groups around the central atom will be most stable when they are as far apart as possible. We call this valence shell electron pair repulsion theory. Because electrons are negatively charged, they should be most stable when they are separated as much as possible. The resulting geometric arrangement will allow us to predict the shapes and bond angles in the molecule.

TO DETERMINE VSEPR SHAPE (electron geometry) 1) Draw the Lewis dot structure for the molecule 2) Identify the central atom 3) Count the number of electron groups around the central atom. 4) Look up the VSEPR shape on the chart. **shapes with no lone pairs are symmetrical **shapes with lone pairs are assymmetrical

Electron Groups Each lone pair of electrons = one electron group on a central atom. Each bond = one electron group on a central atom (regardless of whether it is single, double, or triple.) O N • There are three electron groups on N One lone pair One single bond One double bond

Two Electron Groups: Linear Electron Geometry When there are two electron groups around the central atom, they will occupy positions on opposite sides of the central atom. This results in the electron groups taking a linear shape.

Linear Geometry

Three Electron Groups: Trigonal Planar Electron Geometry When there are three electron groups around the central atom, they will occupy positions in the shape of a triangle around the central atom. This results in the electron groups taking a trigonal planar shape.

Trigonal Planar Geometry

Four Electron Groups: Tetrahedral Electron Geometry When there are four electron groups around the central atom, they will occupy positions in the shape of a tetrahedron around the central atom. This results in the electron groups taking a tetrahedral shape.

Tetrahedral Geometry

Molecular Geometry The actual geometry of the molecule may be different from the VSEPR shape. Lone pairs repel bonded atoms which distorts the expected shape.

Bond Angle Distortion from Lone Pairs Electron Geometry: Tetrahedral Tetrahedral Tetrahedral Molecular Geometry: Tetrahedral Trigonal Pyramidal Bent

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Predicting the Shapes around Central Atoms Draw the Lewis structure. Determine the number of electron groups around the central atom. Classify each electron group as a bonding or lone pair, and count each type. Remember, multiple bonds count as one group. Look it up

Practice: Determine the shape. 1. NF3 2. SiCl4 3. H2O

Types of Bonds: Nonpolar covalent equal sharing of electrons between atoms; occurs between the atoms in a diatomic molecule (HOFBrINCl) and between C and H; ex. CH4 Video is 10:24

Polar Covalent unequal sharing of electrons between atoms; occurs between two nonmetals or a nonmetal and a metalloid; ex. H2O Electrons

Ionic complete transfer of electrons occurs between m/nm (ex. NaCl) m/PAI PAI/nm PAI/PAI

This is a continuum. It describes the “ionic character” of the bond.   Bond type Non-Polar Covalent Polar Covalent Ionic NPC PC I Difference in electronegativity values Distance between atoms on the periodic table Small medium big This is a continuum. It describes the “ionic character” of the bond.

Practice: What type of bond exists in each of the following? 1. HCl 2. CaO 3. H2O Br2

Water is a POLAR molecule The more electronegative atom will have a slight negative charge, the area around the least electronegative atom will have a slight positive charge.

Symmetric molecules tend to be nonpolar Asymmetric molecules with polar bonds are polar Symmetric means there are no lone pair around the central atom.

Bonding determines some physical properties Type of Compound Elements involve in bonding (metal/non metal) Valence electrons are… Melting /Boiling point Electrical conductivity Other properties Metallic Metal-metal delocalized high Conductor Malleable, ductile, shiny Ionic Metal - nonmetal Lost/gained Conducts in solutions or molten Brittle, solid at room temperature Covalent Nonmetal - nonmetal shared low Non-conductor Mostly liquid or gas at room temperature Type of Compound Elements involve in bonding (metal/non metal) Valence electrons are… Melting /Boiling point Electrical conductivity Other properties Metallic Ionic Covalent