Chemistry Exam Study Guide PDF

1.1 Standard atomic notation Atomic number -​ The number of protons in an atom of a certain, unit Mass number -​ The number of protons and neutron in an atom’s nucleus Atom mass unit -​ 1/12 the mass of one atom of carbon -12, unit symbol “u” -​ 1u = 1.660 540 2 x 10^-27 1.2 Ions and o...

1.1 Standard atomic notation Atomic number -​ The number of protons in an atom of a certain, unit Mass number -​ The number of protons and neutron in an atom’s nucleus Atom mass unit -​ 1/12 the mass of one atom of carbon -12, unit symbol “u” -​ 1u = 1.660 540 2 x 10^-27 1.2 Ions and octet rule The octet rule -​ Elements with a full valence shell have special stability, known as a full or stable octet -​ Atoms that do not have a stable octet, tend to combine with other atoms of attain this electron arrangement (this is known as the “octet rule”) -​ This can happen in 3 different ways: Share, lose of gain electrons -​ Any atom losing or gaining electrons will form an ion; an entity with a positive or negative charge Formation of ions -​ Depending on the element, atoms tend to either gain electrons, or will lose electrons to form ions Positive ions: cations -​ Metals tend to lose their few electrons to form a positive ion (cation) Negative ions: anions -​ Nonmetals tend to gian electrons to form a negative ion -​ Naming of anions (replace the end of element name with suffix “Ide” -​ Ex. Chloride Ion Multivalent elements -​ Elements that can form two or more different stable ions -​ Most transition metals are multivalent Polyatomic Ions -​ An ion compound of more than one atom that acts as a unit NICK the CAMEL ate a CLAM for SUPPER in PHOENIX -​ The mnemonic can helps with memorizing the “parent” polyatomic ions that end in “ate” -​ First letter represents the element that is not oxygen -​ The number of consonants = number of oxygen atoms -​ The number of vowels = charge -​ “Ate” suffix = more oxygen atoms -​ “Ite suffix = one less oxygen than “ate” Ions in the human body -​ About 99% of the body is made up of 6 elements: Oxygen, hydrogen, carbon, nitrogen, calcium and phosphorus -​ Smaller quantities of sulfur, chlorine, sodium, magnesium, iodine, and iron, and they play key roles in the body (many are dissolved in water) 1.3 Isotopes, radioisotopes and atomic mass Isotopes -​ A different form of the same element, same number of protons, but different number of neutrons -​ Mass number is used to distinguish between different isotope -​ Isotopes exist in different relative abundances (ex. 78.7% Mg-24, 10.1% Mg-25 and 11.2% Mg-26) -​ The percentage of isotope in a sample is known as isotopic abundance Radiation and radioisotopes Radioactive -​ Potential to emit radiation upon decay Radioactive decay -​ Spontaneous disintegration of unstable isotopes Radioisotope -​ Isotope that spontaneously decays to produce two or more small nuclei and radiation Nuclear radiation is very small particles released from the nucleus of radioisotopes, such as: -​ Alpha particle: positively charged particle with same structure as nucleus of helium atom -​ Beta particle: a negatively charged particle indentical to an electron -​ Gamma ray: high energy electromagnetic radiation What is this mass…?? Mass number -​ Number of protons and neutrons Average atomic mass -​ The average mass of all the different isotopes of an element in a specific sample Relative atomic mass -​ Average of the weight of all isotopes in a normal environment on the Earth’s crust, standardized number determined by IUPAC Calculating average atomic mass -​ Isotopes are not found in even amounts in nature, some are more common than others. Therefore, your “average” mass of all isotopes of an element can be calculated 1.4 The periodic table and periodic law Reviewing the table -​ Groups and chemical families -​ Periods -​ Metals -​ non-metals Groups or chemical families Vertical columns in the periodic table -​ 1A or group 1 are called alkali metals because they react with water to form an alkaline solution (soft, silvery) -​ Group IIA or 2 are called the alkaline earth metals because they are reactive, nut not as reactive as Group 1A (They are soft metals but less reactive than group 1) -​ Group VIIA or 17 are the halogens (They need only one electron to fill their outer shell and they are very reactive) -​ Group VIIIA or 18 are the noble gasses as they have completely filled outer shells, they are almost non reactive “inert” (also tasteless, odorless, and colorless) Transition metal -​ Elements in groups 3-12 -​ Less reactive, harder metals, variable reactivity -​ Includes metals used in jewelry and construction -​ Metals used “as metal” Inner transition metals -​ Lanthanides: includes elements with atomic number 57 to 70 -​ Actinides: includes elements with atomic number 89 to 102 -​ Transuranic elements: synthetic elements with atomic numbers 93 or greater Factors influencing trends in the periodic table Number of protons -​ Effective nuclear charge (attractive pull of the positive nucleus on the electrons) -​ More protons = more attractive force Number of orbits  -​ Electron shielding (inner core electrons shield the outer valence electrons from the attractive force of the nucleus -​ More orbits = less attractive force Trends in the periodic table -​ Elements in the same group have similar chemical and physical properties -​ Hydrogen typically behaves like a non-metal, but under high pressures and low temperatures, it behaves like a metal -​ Reactivity of metals (increases towards left) (increases going down) -​ Reactivity of non-metals (increases towards right) (increases going up) Atomic radius -​ A measurement of the distance between the nucleus to just beyond the outermost electron -​ More orbits = larger atom -​ More protons = smaller atom (pill e- in closer) Ionic radius -​ Is the radius of an ion larger or smaller than that of the atomic radius of the same element? -​ It depends on whether or not it gains or loses electrons -​ Cations (+) are always smaller -​ Anions (-) are always larger ​ Bohr-rutherford diagrams Lewis dot diagram -​ Shows only valence electrons -​ Add electrons clockwise, starting from the top before beginning to pair them 1.5 Ionization energy and electron affinity Ionization energy Valence electrons are bound to an atom by their attractive force to the nucleus -​ Removing electrons requires energy -​ Adding electrons releases energy Ionization energy (creates plasma) -​ The amount of energy needed to remove an electrons from an atom or ion in the gaseous state -​ Unit for ionization energy is KJ/mol -​ First ionization energy = energy required to remove the most loosely held electron -​ Second ionization energy = energy required to remove the next most loosely held electron -​ IE1 < IE2 < IR3 Ionization energy periodic trends -​ Less energy is required as you move down a group because the further the electron is from the nucleus, the easier it is to remove -​ More energy is required as you move from left to right across a period because the electrons are closer to the nucleus -​ Larger atom = easier to remove Electron affinity -​ The energy change that occurs when an electron is added to a neutral atom in the gaseous state  -​ Electron affinity for helium is below zero, meaning it actuall requires energy for an electron to be added because the repulsive force is greater than the attractive forces Electronegativity -​ Ability of an atom to attract a bonding electron to itself -​ Same forces and factors come into play for electronegativity as IE and EA  -​ Distance from nucleus: smaller the atom, closer the nucleus is to neighboring electrons = stronger attraction -​ Effective nuclear charge: large Zeff’ greater the attraction of the nucleus -​ Pauling scale measures relative electronegativity 1.6 Electronegativity & Bond Polarity Electronegativity -​ The ability of an atom to attract bonding electrons to itself -​ The higher the electronegativity, the better it is a attracting electrons -​ Generally, electronegativity increases from left to right and from bottom to top -​ Fluorine has the highest electronegativity (4.0) and Francium has the lowest (0.7) Electronegativity difference (∆EN)  -​ The difference in electronegativity between two bonded atoms or ions -​ Helps to predict whether a bond is ionic or covalent -​ The greater the difference, the more likely electrons are to be transferred (ionic bond) E.g. ∆EN between Na (3.2) and CL (0.9) 3.2-0.9 = 2.3 If ∆EN is > 1.7 then the bond is considered ionic If ∆EN is ≤ 1.7 then the bond is considered covalent Polar and nonpolar bonds Non-polar covalent bonds -​ Bonds between atoms with same (or very similar) electronegativity -​ Atoms share electrons equally, ∆EN=0