Atom nucleus (protons and neutrons) electron cloud (electrons)

Transcription

1 Atom nucleus (protons and neutrons) electron cloud (electrons) Atomic Number equal to the number of protons Mass Number protons + neutrons Charge when # of electrons # of protons Negatively Charged Ion (anion) has more electrons than protons Positively Charged Ion (cation) has more protons than electrons Neutral Atom - # of electrons = # of protons All of this information may be written as part of a chemical symbol: mass number 15 7 atomic number 3 N charge

2 Isotopes So far, you know that the mass number is equal to the number of neutrons plus the number of protons. This number is important in our discussion of elements (and their atoms), but why? Let s investigate In nature, elements are made up of atoms that are not exactly the same some are heavier than others but why?? Consider the relative masses of the sub-atomic particles: Protons atomic mass units (amu) Neutrons amu Electrons amu Since the number of protons NEVER changes, they are not responsible for the changes in mass. Since electrons weigh pretty much NOTHING compared to the other particles, they are not responsible for the changes in mass. Therefore, it is the change in the number of neutrons that is responsible for the changes in mass.

3 In nature, elements are made up of atoms that have slightly different numbers of neutrons. This results in the different varieties of atoms called isotopes. Thus, it is important to understand how to interpret isotopic symbols; it allows us to figure out WHICH version of the element (atom) we are dealing with. So For an element, the atomic number (number of protons) NEVER changes, but the mass number will change depending on the isotope. Some additional info These different versions of atoms are taken into account when determining the atomic mass. On the periodic table, the atomic mass is the average of all the masses of the isotopes. Example: C-12 and C-14

4

5 Lewis (Electron) Dot Diagrams Atoms tend to gain or lose electrons in a n attempt to satisfy the octet rule. That is, to become like the noble gases, which are the most stable elements. Atoms can share or give/take electrons. Electrons are being used to form bonds. In order to draw atoms and their electrons in the most concise and efficient manner, Lewis suggested that we only include the valence electrons, since they are the only ones that are involved in the bonding. Examples: Na Be Al Sn N S F Ar

6 Give/Take Diagrams used to show the transfer of electrons (ionic bonding ) Examples: Na Cl NaCl (sodium chloride) Li O Li 2 O (lithium oxide) Li Be F F BeF 2 (beryllium fluoride) Question: What do you notice about how the electrons move? Electrons are given from the metals Electrons are taken from the non-metals Electrons are always transferred from the atom with the least electrons to the ones with more (it is easier)

7 How to Draw Lewis Structures Lewis structures are a way to write chemical compounds where all the atoms and electrons are shown. Steps to follow: 1. Count the total valence electrons for the molecule: To do this, find the number of valence electrons for each atom in the molecule, and add them up. 2. Figure out how many octet electrons the molecule should have, using the octet rule: The octet rule tells us that all atoms want eight valence electrons (except for hydrogen which only wants 2), so they can be like the nearest inert gas. Use the octet rule to figure out how many electrons each atom in the molecule should have, and add them up. 3. Subtract the valence electrons from octet electrons: Or, in other words, subtract the number you found in step 1 from the number you found in step 2 above. The answer you get will be equal to the number of bonding electrons in the molecule. 4. Divide the number of bonding electrons by two: Remember, because every bond has two electrons, the number of bonds in the molecule will be equal to the number of bonding electrons divided by two. 5. Draw an arrangement of the atoms for the molecule that contains the number of bonds you found in step 4 above: Some handy rules are: - hydrogen always has one bond, never more. - the halogens usually have one bond. - the family that oxygen is in usually makes two bonds. - the family that nitrogen is in usually makes three bonds. - the family that carbon is in always makes four bonds. 6. Find the number of lone pair (nonbonding) electrons by subtracting the bonding electrons (step 3) from the valence electrons (step 1). Arrange these around the atoms until all of them satisfy the octet rule: Remember, ALL elements EXCEPT hydrogen want eight electrons around them, total. Hydrogen only wants two electrons.

8 CHEMICAL BONDING - chemical bonds are attractive forces. - valence electrons are the only electrons that take part in chemical bonds. They are the outer energy level electrons so they are available. There are two different types of chemical bonds that occur between atoms: ionic bonding and covalent bonding. Note: i) chemical bonds involve a competition for bonding electrons in unfilled orbital. ii) bonds continue to form until each atom has filled its valence orbital (octet rule).

9 IONIC BONDING - ions that are formed when neutral atoms either gain or lose valence electrons. ion - an atom that has gained or lost electrons. Cation electron neutral electron Anion + loss atom gain - - the driving force behind the formation of ions is the tendency of atoms to reach a more stable electron configuration (to have a stable octet). This is done by acquiring either a completely full or a completely empty valence level. This gives the ion the same stable electron configuration as a noble gas. Example. 1. S + 2 e - = [ S ] 2-2. Na - 1 e - = [ Na ] +

10 - the ions with opposite charges will attract one another by electrostatic forces. These forces of attraction that bind oppositely charge ions are called ionic bonds. ATOMS IONS IONIC COMPOUND 1. Na + Cl = [ Na ] + + [ Cl ] - = NaCl 2. Ca + 2 Cl = [ Ca ] [ Cl ] - = CaCl Al + 3 O = 2 [ Al ] [ O ] 2- = Al 2 O 3 Note: 1. The number of ions in the formula is adjusted so that each molecule is neutral overall. 2. Each element in the compound has achieved a noble gas configuration.

11 Use electron dot diagrams to show the ionic bonding process with these elements. 1. Potassium + Oxygen 2. Magnesium + Nitrogen 3. Aluminum + Bromine 4. Potassium + Iodine 5. Calcium + Sulfur 6. Sodium + Phosphorus 7. Aluminum + Sulfur 8. Sodium + Oxygen

12 COVALENT BONDING - some atoms share electrons in order to attain stable octets. By sharing electrons, two fluorine atoms can complete their octets and become more stable. This process can be shown with electron dot diagrams. When one pair of electrons is shared between two atoms, it is called a single covalent bond. F + F = F F - atoms can share more than one pair of electrons. Double covalent bonds involve two shared pairs. O + O = O O - nitrogen atoms combine with a triple covalent bond. N + N = N N Diatomic Molecules: molecules that contain two of the same atoms. Example: Iodine, bromine, chlorine, fluorine, oxygen, nitrogen, hydrogen

13 POLYATOMIC IONS tightly bound groups of atoms that behave as a unit and carry a charge. there are many of these ions. Since we cannot predict what their charges will be, you will be given a list to work with. when there is a compound formed with polyatomic ions, both ionic and covalent bonding takes place.

14 Properties of Ionic and Covalent Bonding Properties of Ionic Compounds: 1. They have relatively high melting points due to strong attraction in bonds 2. They conduct electricity when molten or dissolved in water (ions can move freely) 3. Those in solid state are not electrical conductors (ions cannot move) 4. Ionic compounds dissolved in water form electrolytic solutions. Electrolyte a substance that dissolves in water to produce a solution that conducts electricity. Properties of Covalent Compounds: 1. They have relatively low melting points due to weaker attraction in bonds 2. They tend not to conduct electricity when in solid or liquid state or dissolved in water (do not form ions; non-electrolytes)