Download Regents Review Sheet1-3

Regents Review Sheet 1 – Atomic Structure
These are some important points to remember about Atomic Structure.
Use this sheet when you do the Review Problems.
Points of Interest
• Law of Conservation of Matter-matter can not be created or destroyed
• Law of Constant Composition-a compound always contains the same
elements in the same proportion or ratio by weight
Atom - The smallest particle of an element that retains the chemical
identity of the element
• A) Electrons
• B) Nucleons-particles in the nucleus; 2 kinds
– 1. Protons-positively charged
• Have a mass of 1 amu (atomic mass unit)
– 2. Neutrons-neutral
• Also have a mass of 1 amu
Experiment
• Rutherford did an experiment where he shot alpha particles (very
small particles) at a piece of gold foil
– He found that most of the alpha particles go right through, but
that occasionally one of the particles was deflected a whole lot,
as if it had ricocheted off of something solid. It hit the nucleus.
Atomic Number-Tells the number of protons in the nucleus
Ions-atom with an electrical charge
– Anions are negative ions (GAINED –e)
– Cations are positive ions (LOST -e)
Isotopes-same elements have different numbers of neutrons
The Bohr Model of the Atom
• The closer the electron is to the nucleus, the lower the energy
• Note: only a certain amount of –e fit into a given energy level
• These principal energy levels approximate how far an electron is from
the nucleus
Ground State-When the electrons are in the lowest available energy levels,
the atom is in the ground state
Excited State - If the atom absorbs energy, the electrons become “excited”
and may jump up to a higher energy level
Orbitals - The space within an atom where an electron or pair of electrons is
likely to be found
Valence Electrons - Electrons in the outermost principal energy level (also
called shell) of an atom
• Cl has 17 electrons. Of these, 7 are in the 3rd principal energy level.
Therefore, it has seven valence electrons
• The kernel is the atom except for the valence electrons
Ionization Energy-The amount of energy required to remove the most
loosely bound electron from an atom in the gaseous phase
Reference Table Information for this Unit
A.
Periodic Table – gives Atomic Mass, Atomic Number, Electron
Configuration, etc.
Regents Review Sheet 2 – Periodic Table
These are some important points to remember about Periodic Table. Use this
sheet when you do the Review Problems.
Points of Interest
• The periodic table organizes the elements by their properties
• The elements are arranged in rows called “periods” and columns
called “groups”
• The groups hold elements that are similar and have related properties
• The periods (7 of them) represent the principal energy levels. The 1st
period contains elements with electrons in the 1st principal energy
level. The 4th principal energy level has elements with electrons in the
4th principal energy level
• The elements are arranged in increasing atomic number (in increasing
numbers of protons in the elements)
• Groups (columns) are labeled 1 -18
• Group 1 is the alkali metals
• Group 2 is the alkaline earth metals
• Group 17 is the halogens
• Group 18 is the noble gases
• In the table, there is usually a “staircase” drawn toward the right-hand
side; the elements to the left are metals, and the elements to the right
are non-metals
• Bordering the steps are the metalloids or semi-metals; B, Si, As, Te,
Ge, Sb, At
• Most elements are solids at room temperature except:
– Hg and Br
liquids
– H, O, N, F, Cl and the noble gases, which are all gases
Trends and Properties
• 1) Atomic Radius- half the distance between the centers of 2 atoms of
the element that are just touching each other
– Atoms get larger as you go down a group; this makes sense
– Atoms get smaller as you go across a period from left to right.
It occurs because as an atom gets more protons, the electrons
are more strongly attracted to the nucleus, and the atoms radius
decreases.
– Also as an electron is gained or lost by an atom, the radius
changes
- If you add electrons, forming a negative ion, you get a
larger radius
- If you lose electrons, forming a positive ion, you get a
smaller radius
• 2) Ionization Energy- The amount of energy required to remove the
most loosely bound electron from an atom in the gaseous phase
• Ionization Energy decreases as you move down a group
• It increases as you go across a period from left to right
These trends occur because larger atoms are not able to hold their
electrons as tightly as smaller atom
• Electronegativity - How well an atom attracts another atom’s
electrons
Regents Review Sheet 3 – Bonding
These are some important points to remember about Bonding. Use this sheet
when you do the Review Problems.
Points of Interest
Ionic Bonds-really an attraction between a positively charged metal ion and a
negatively charged non metal ion
– During a reaction between these metals and the non-metals, electrons in
the valence shell of the metal are transferred to the valence shell of the
non- metal
– Note: By forming ionic bonds, a metal and non- metal achieve the stable
octet in the valence shell
Covalent bonds are formed by a shared pair of electrons between 2 atoms in a
molecule
- Bonds-the atoms share 4 –e (2 pairs of –e)
-Triple Bond-the atoms share 6 –e (3 pairs of –e)
•
•
General Rules for Lewis Dot Diagrams
1) On a sheet of paper write the element symbol with the number of valence
electrons around it
2) Put all the atoms together in such a way as to satisfy the octet rule
Octet Rule
• Atoms tend to gain, lose or share electrons in order to acquire a full set of valence
electrons
• In other words, each atom will become involved in bonds which will result in
them becoming like noble gases
Crisscross Method
• If you have two atoms that you know will react, there is an easy way to see what
the molecular formula of the compound they will form is.
– 1) Look at the periodic table and find the charges in the upper right hand
corner of the box. Write these numbers to the upper right of the atomic
symbol
– 2) Crisscross the numbers
• Example - Find the molecular formula of the product of a reaction between Ca
and F
– 1)
Ca2+ F1– 2)
Ca1 F2
ElectronegativityShows the strength of attraction that one atom has for an electron
•
•
Different atoms have different electronegativities; what this means is that
some atoms attract electrons better than other atoms do
The atom with the stronger electronegativity will actually hold the shared
electron closer to itself than to the other atom in the covalent bond
•
Electronegativity Difference
• Shows whether a bond is Ionic, Polar Covalent or Non-Polar Covalent
• If the difference is ≤ 0.4 it is Non-polar Covalent
• If the difference is between 0.4 and 1.7 it is Polar Covalent
• If the difference is > 1.7 the bond is Ionic
You can also determine if a bond is Ionic, Polar Covalent or Non-Polar Covalent in this
way
Ionic – between a metal and a non- metal
Polar Covalent – between 2 different non- metals
Non-Polar Covalent – between 2 identical non- metals (diatomic molecules)
Coordinate Covalent Bond - When one atom in the covalent bond has donated both of the
electrons in the shared pair – (example is Ammonia)
Network Solids - Atoms are linked throughout the sample by covalent bonds; this makes
the substance very hard and strong
• The substance will also have very high melting points
• They are poor conductors of heat and electricity
• Diamonds, silicon carbide and silicon dioxide are examples
Metallic Bonding - Occurs between atoms of metals
• Metals form + ions; these + ions are linked by an attraction to a bunch of free
floating –e (this is why metals conduct electricity)
• The metal ions don’t have full valence shells, so they pick up –e and lose –e often
• This makes these substances very malleable
Molecular Attraction or Interactions
• Remember polar covalent bonds-molecules that have this type of bond are called
dipoles
• Dipoles- one side of the molecule is slightly positive and one side is slightly
negative.
• Example is HCl - each HCl has a + side and a – side. When 2 HCl molecules
come in contact with each other, the + H of one becomes attracted to the – Cl of
the other and they stick to each other
Hydrogen Bonds
• Occur when H is bonded to an element with a small atomic radius and high
electronegativity
• H becomes slightly + since the other element doesn’t share the electron evenly
with it
•
Because of this slight + charge, H is attracted to - atoms
Attractions and Molecule Shape
• When you have more than 2 atoms in a molecule, the shape of the molecule
determines its polarity; when polar bonds are uniformly distributed around a
molecule, the slight charges cancel each other out and the molecule is actually
non-polar
Van Der Waals Forces
• Even without polar attractions and H-bonds, there are still weak attractive forces
between molecules
• A trick to remembering the characteristics of Van Der Waals forces is to think of
it as being like gravity; large atoms have powerful VDW forces, while small
atoms have weak VDW forces.
• Also, the closer the atoms the more powerful the VDW forces.