Unit 9 Bonding Test Wed. 2/17

Why do atoms form bonds? to have the most stable electron configuration for its electrons to achieve a lower, more stable, energy state 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.

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

Valence electrons To find the of valence electrons in an atom of a representative element, look at the last digit of its group number.

Electron dot structure Also called “Lewis valence electron dot structure”

Octet rule In forming compounds, atoms tend to achieve the electron configuration of a noble gas Fill up their highest energy level with 8 electrons OR Empty their highest energy level to Zero electrons

cation I’m “paws-itive” A positively charged ion, or cation, is produced when an atom loses one or more valence electrons.

Formation of a cation

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anion – formed when an atom gains valence electrons

Nonmetals attain noble-gas configurations by gaining electrons.

“losers on the left” GAIN VALENCE ELECTRONS LOSE VALENCE ELECTRONS

PRACTICE How many valence electrons? Potassium Carbon Magnesium Oxygen

Ionic compound Composed of cations and anions Ionic compounds are electrically neutral.

Ionic bonding video 1 Ionic bonding video 2 IONIC BOND an electrostatic* force that holds ions together in ionic compounds *Electrostatic refers to the attraction between opposite charges Ionic bonding video 1 Ionic bonding video 2 Ionic Bond

Chemical formula Shows the number of each element in the smallest representative unit of a substance.

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. 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

Writing Lewis Dot Structures – Ionic Bonds Write the Lewis dot structure for each type of atom involved in the compound. Determine which atom will lose electrons (metal) and which atom will gain electrons (nonmetal). Draw arrows to show the electrons moving from the metal to the nonmetal. If an octet is not filled for the nonmetal or if you don’t use up all of the metal’s valence electrons, you will need add more nonmetal or metal atoms. Draw as many atoms of the metal as necessary to fulfill the octet for the nonmetal and use all of the valence electrons from the metal. Consolidate your diagram. Draw the metal on the left with its charge in the upper right corner. (If more than one of these was needed, write a coefficient in front of the metal.) Draw the nonmetal with its octet. Enclose it in brackets. Include the charge outside of the brackets in the upper right corner. (If more than one of these was needed, write a coefficient in front of the nonmetal.)

Properties of ionic compounds (all related to the strong attraction between the + and – charges) hard, crystalline solids at room temperature high melting points (and boiling points) good conductors – in aqueous solutions (dissolved in water) and when molten (melted)

not crystalline

Metallic bonding To describe the nature of metallic bonding, consider the valence electrons of atoms in a pure metal to behave as a 'sea' of delocalized electrons. 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 bond the forces of attraction between free-floating (delocalized) electrons and positively charged metal ions. Watch This 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

Ductility – property of a metal that enables it to be drawn into a wire Malleability- means that a metal can be hammered into a sheet

Electrical conductivity Thermal conductivity To conduct an electric current through the flow of electrons Transfer heat

video1 Video 2

alloy A mixture of 2 or more metals (KP p. 224) Alloys are important because their properties are often superior to those of their component elements.

Really, we don’t hate you.

Covalent Bond formed when atoms are held together by sharing electrons 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.

Number of bonds single covalent bond - when 1 pair of electrons is shared between 2 atoms double covalent bond – when 2 atoms share 2 pairs of valence electrons; ex. O2 Triple covalent bond – when 2 atoms share 3 pairs of valence electrons; ex. N2

(Review) Diatomic Molecules HOFBrINCl Share electrons when they bond together

Polyatomic Ions covalently bonded group of atoms, with a + or - charge Watch this

Bond type video Transfer of electrons between atoms occurs between m-nm unequal sharing of electrons between atoms occurs between other combinations of nm-nm or nm-metalloid equal sharing of electrons between atoms occurs between the two atoms in a diatomic molecule (H-H) and between C and H atoms (CH4)

Types of Bonds: Video is 10:24

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.

How to draw Lewis dot structures for covalent molecules. 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)

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.

valence shell electron pair repulsion VSEPR Theory also called electron geometry Electron groups around the central atom will be most stable when they are as far apart as possible.

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. -an electron group can be a lone pair or a bond (single, double, or triple) 4) Look up the VSEPR shape on the chart. **shapes with no lone pairs are symmetrical **shapes with lone pairs are assymmetrical

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

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 (no lone pairs around the central atom) molecules tend to be nonpolar. Asymmetric (at least one lone pair around the central atom) molecules with polar bonds are polar.

CHARACTERISTICS OF IONIC AND MOLECULAR COMPOUNDS Video CHARACTERISTICS OF IONIC AND MOLECULAR COMPOUNDS CHARACTERISTIC IONIC COMPOUND MOLECULAR COMPOUND Representative unit   Bond Formation Type of elements Physical state Melting point Solubility in water Electrical conductivity of aqueous solution