Chapter 6 & Chemistry Review Chemistry Unit 4 VSEPR Theory Valence Shell Electron Pair Repulsion ● If electrons repulse one another, then an electron pair also repulses electron pairs. ● Molecular shape assumes the geometry that minimizes electrostatic repulsion ○ Electron pairs must be as far apart as possible Simplified version of molecular geometry 3 atoms – BeF2 – linear 4 atoms – BF3 –no lone pair e-Trigonal Planar 4 atoms –NH3 – A lone pair e-Trigonal pyramdical Electron affinity ● Ability of an atom to accept an electron ● Do the noble gasses (halogens) have a strong electron affinity? ● No “8” electrons Neutral atom has no positive or negative charge. When an atom contains an equal number of electrons and protons, then the atom will be neutral. If the number of electrons becomes more or less than neutrons then the atom acquires charge and becomes an ion. Electron affinity is the energy change that occurs when an electron is acquired by a neutral atom or attracts an electron. Electronegativity In the Pauline Scale, Fluorine has the highest electronegativity of 4.0 What is the force that holds two atoms together? A bond What is an ionic bond? A bond between a metal and non-metal What is a covalent bond? A bond between two non-metals Factors Affecting Electron Affinity: •Nuclear Charge: •A higher nuclear charge leads to a stronger attraction for the added electron, and thus a higher electron affinity. •Atomic Radius: •A smaller atomic radius leads to a stronger attraction for the added electron, and thus a higher electron affinity. •Electron Shielding: •The core electrons shield the valence electrons from the full nuclear charge, reducing the attraction for the added electron. •Electron Configuration: •Elements with stable electron configurations (like noble gases) tend to have low electron affinities. Ionic Bonding ● Hy Ionic Bonding Types of Covalent Bonds Double Bond Single Bond ● ● ● ● One pair of electrons is shared between two atoms -total 2 e Weakest covalent bond to form Easiest covalent bond to break Ex: H2 CH4 H2O ● ● ● ● Two pairs of electrons are shared between two atoms total 4e Medium covalent bond to form Medium covalent bond to break Ex: C2H4 CO2 Triple Bond ● ● ● ● Three pairs of electrons are shared between two atoms total 6e Strongest covalent bond to form Most difficult covalent bond to break Ex. N2 CO CN- How can you tell if a bond is ionic or covalent? ● Look at their location on the Periodic table. ● If one is a metal – ionic ● Both have to be a non-metal - covalent Intermolecular Forces 1. Dipole-Dipole a. Hydrogen Bonding 2. London Dispersion Forces Intermolecular Forces Di – two Dipole – two poles + & - Dipole-dipole are from the constant motion of the Intermolecular Forces LDF electrons stay stationery The force that holds them together is the negative electrons to the positive proton Hydrogen Bonds • Are bonds that include hydrogen atom • Not a sharing or transfer of electrons, but an attraction with the hydrogen to another atom • Dipole to Dipole attraction between molecules More Bonds Polar Bonds ● Unequal sharing of electrons ● Geometry can affect polarity ● Different atoms bonded will generally be polar ● Asymmetrical structure is always polar Nonpolar Bonds ● Equal sharing of electrons ● Like atoms will always be nonpolar ● Symmetrical structure is always nonpolar ● Resonant Structure: Must contain a double bond The double bond can go between any pair Lewis Dot Diagrams By knowing the valence electrons, a Lewis Dot diagram can be made Lewis Dot Diagram or Lewis Structure is a visual representation of the number of valence electrons in an atom or ion. Why would we only need to Show the valence? electrons? Putting it all Together Lewis Dot Arrange ➢ Draw Lewis Dot structure of each atom ○ Add up all valence electrons for all atoms ➢ Put the least electronegative atoms in the center. ○ Farthest from F on PT ○ Listed first in formula ○ Typically be the atom there are the fewest of ○ Can form the most bonds ○ Never be H ➢ Place the other atoms around the central atom symmetrically Connect ➢ Use a dash — or two dots in between atoms to represent a single bond ➢ Give all atoms (except H) an octet ➢ Count your valence electrons used (should be the same as step one!) ○ If there are extra e-, put lone pairs on the central atom Polar molecules vs Non-Polar Molecules IF they are symmetrical on all sides – Non-polar. IF they are not symmetrical, then they are polar Draw Lewis dot structures • • • Add up all the valance electrons Put electron pairs about each atom such that there are 8 electrons around each atom (octet rule), with the exception of H, which is only surrounded by 2 electrons. • • O2 Atomic # O - 8 2s 2s 4p – 6 valence e Valence electrons = the column # Draw Lewis dot structures • • • Add up all the valance electrons Put electron pairs about each atom such that there are 8 electrons around each atom (octet rule), with the exception of H, which is only surrounded by 2 electrons. • • • O2 Atomic # O - 8 2s 2s 4p – 6 valence e per O 12 Valence e Draw Lewis dot structures • • • • Add up all the valance electrons Put electron pairs about each atom such that there are 8 electrons around each atom (octet rule), with the exception of H, which is only surrounded by 2 electrons. Each pair of dots = a “—” • 12 Valence e Remember the rule of 8 Draw Lewis dot structures • • • • • Add up all the valance electrons Put electron pairs about each atom such that there are 8 electrons around each atom (octet rule), Least electronegativity in the center (lower right of the Periodic Table) or most willing to share e • • • H2O Atomic # O – 8 6 val e Atomic # H – 1 1 val e x 2 8 valence e Draw each atom separately then combine to make 8 e Draw Lewis dot structures • • • • Add up all the valance electrons Put electron pairs about each atom such that there are 8 electrons around each atom (octet rule) • Carbon C and Chlorine Cl C 4 val e Cl 7 val e Draw Lewis dot structures • • Add up all the valance electrons Put electron pairs about each atom such that there are 8 electrons around each atom (octet rule), with the exception of H, which is only surrounded by 2 electrons.
