sci 10

• Big Bang and Stellar Nucleosynthesis paved the way to create elements and these elements compounds.
• Atom - the smallest particle of an element that can exist and still have the properties of the element.
• Compound - pure substances that are composed of elements in definite positions
• Dalton’s Atomic Theory - 1808, English school teacher and chemist.

  • Each element is made up of tiny particles called atoms.
  • The atoms of a given element are identical.
  • Chemical reactions involve reorganization of the atoms—changes in the way they are bound together. 
  • The atoms themselves are not changed in a chemical reaction.
• J.J Thomson

  • 1897, British Physicist
  • Discovered the electron in his cathode ray tube. 
• E. Rutherford (1908) - Radon and some elements emit streams of positively charged particles (alpha particles).

  • Gold Foil Experiment
• 20th Century - The development of mass spectrophotometers led to the discovery of the neutrons
• Atomic Radius

  1. Groups

  • Atomic radius increases as you move down a group because there are more electrons in more Principal Energy Levels.
  • Atomic size increases

  2. Periods

  • Atomic radius decreases as you move across a period, why —  (-) electrons increase, but so do (+) protons !!! 
  • Increased (+) nuclear charge pulls the (-) electrons closer to the nucleus.
  • Atomic size decreases.
• Ionization Energy

  • The energy needed to overcome the attraction of the nuclear charge and remove an electron (from a gaseous atom).
  • Energy is needed to remove an electron from an atom.
• Electronegativity

  • Ability of an atom to attract electrons towards itself in a chemical bond.
• Electronegativity — Periods

  • Electronegativity increases going left to right across the periodic table.
• Electronegativity — Groups

  • Electronegativity decreases going down a group.
  • The bonding electrons are increasingly distant from the attraction of the nucleus.
• Electron Affinity

  •  A measure of the energy absorbed when an electron is added to a neutral atom to form a negative ion.
  •  Most elements have a negative electron affinity. This means they do not require energy to gain an electron; instead, they release energy.
• Nonmetallic Property

  • decreases within a group and increases within a period 
• Metallic Property

  • increases within a group and decreases within a period
• Element

  • a substance where its atoms have the same number of protons in it
• Compound

  • combination of two or more elements
• The Chemical Bonds

  • attractive force that holds 2 atoms together in a more complex unit
  • form as a result of interactions between electrons found in the combining atoms
• 1. Valence Electron

  -  an electron in the outermost electron shell of a representative element or noble-gas element.
• 1. Lewis Symbol

  • chemical symbol of an element surrounded by dots equal in number to the number of valence electrons present in atoms of the element.
• 1. Octet Rule

  • certain arrangements of valence electrons are more stable than others
  • The valence electrons configurations of the noble gasses are considered the most
  • stable of all valence electrons configurations.
  • to form compounds, atoms of elements lose, gain, or share electrons in such a way as to produce a noble-gas electron configuration for each of the atoms involved
• Types of Chemical Bonds

  1. Ionic

  • chemical bond formed through the transfer of one or more electrons from one atom or group of atoms to another atom or group of atoms.
  • ionic compound a compound in which ionic bonds are present
• Ionic compounds are always neutral; no net charge is present. The ratio in which positive and negative ions combine is the ratio that achieves charge neutrality for the resulting compound.
• 1. Covalent

  • chemical bond formed through the sharing of one or more pairs of electrons between two atoms.
• Molecular Compound (Covalent Compound)

  • a compound in which atoms are joined through covalent bonds
• Valence Shell Electron Pair Repulsion Theory 

  • a model used to predict the shapes of compounds
• Molecular Shapes 

  • Lewis structures give atomic connectivity: they tell us which atoms are physically connected to which.
  • The shape of a molecule is determined by its bond angles.
  • Consider CCl4: experimentally we find all Cl-C-Cl bond angles are 109.5 degrees.
• In order to predict molecular shape, we assume the valence electrons repel each other. Therefore, the molecules adopt whichever 3D geometry minimizes the repulsion. We call this process Valence Shell Electron Pair Repulsion (VSEPR) theory.
• POLAR AND NONPOLAR COVALENT BONDS

  • Polarity means having dipoles, a positive and a negative end.
  • Polar molecules have dipoles. Their dipole moments do not add up to zero (or do not cancel out). Water and carbon monoxide are examples of polar molecules.
• Electronegativity Differences

  > 2.0 = Ionic

  <0.4 = Nonpolar Covalent

  0.5-1.7 = Polar Covalent
• Nonpolar molecules don’t have positive or negative ends. Carbon tetrachloride and methane are examples of nonpolar molecules.
• Generally, you can tell if a molecule is polar or nonpolar based on:

  1. its structure or shape
  2. the polarity of the individual bonds present in the molecule
• CO2 as Nonpolar Molecule

  Exception to the Rule

  1. NaBR
  2. HF
  3. NaCl
• The Molecular Geometry of Water

  Water is bent and is a polar molecule.
• Forces of Attraction in Water

  Water as a Universal Solvent.

  • Solubility (“like dissolves like”)
  • Polar solutes dissolve in polar solvents
  • Sucrose is soluble in water because both of them are polar molecules.
• Properties of Water vs. Other Compound

  water - melting point - 0c

  - boiling point - 100c

  carbon dioxide

  -melting point - -56.6

  boiling point - -78.5
• Types of IMFA

  1. London dispersion forces (LDF)

  • present in all molecules 
  • caused by fluctuations in the electron distribution within atoms or molecules
  • weak type of IMFA
• Intermolecular Forces of Attraction (IMFA)

  • IMFA are the attractive forces present between molecules. Generally, they are called van der Waals forces, named after Johannes van der Waals.
• Types of IMFA

  2. Dipole-Dipole Forces

  • attractive forces between polar molecules
  • result of the electrical interactions among dipoles on neighboring molecules
  • moderately strong type of IMFA
• Types of IMFA Hydrogen Bonding

  • a special kind of dipole-dipole forces
  • an attractive force between a hydrogen atom of one