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Unit 3 – Lessons 1-4 “ATOMS and the PERIODIC TABLE” (Pages 153-208) I. LESSON 1 A. The Atom 1. The basic UNIT of all MATTER is the ATOM 2. ATOM is the SMALLEST particle of an ELEMENT retaining all the chemical PROPERTIES of that ELEMENT 3. Democritus’s GREEK term “ATOMOS,” which means “not able to be divided,” is the ORIGIN of the word, ATOM *4. Individual ATOMS are so SMALL they can only be seen using a scanning tunneling MICROSCOPE (STM), a special type of ELECTRON microscope *a. Due to their SMALL size, huge numbers of ATOMS are used in the composition of very small substances *b. The average DIAMETER (distance across) of an ATOM is about 0.000000002 m or 2 x 10-9 m (scientific notation) or [2 nm (nanometers)] *5. Scientific notation is a system of writing NUMBERS with a LARGE number of place value positions containing zeroes *a. SCIENTIFIC NOTATION is a NUMBER written as a product with two FACTORS 1. The 1st FACTOR is a single digit whole number greater than 0, but less than 10 2. The 2nd FACTOR is a POWER of 10 in exponential form [(e.g.) “105” ] *b. Steps for writing a NUMBER in a SCIENTIFIC NOTATION format: 1. Move the decimal POINT until the first NUMBER is between “1-10” but not 10 2. Count the number of PLACE VALUE positions moved and express that total as an EXPONENT a. When you move the DECIMAL point to the left, the EXPONENT will be expressed as a positive number b. When you move the DECIMAL point to the RIGHT, the EXPONENT will be expressed as a NEGATIVE number 3.(e.g.) 100,000 = 1 x 105 Speed of light 30, 000,000,000 cm/sec = 3 x 1010 cm/sec 150,000,000 = 1.5 x 108 0.0000004 = 4 x 10 -7 0.00000634 = 6.34 x 10 -6 203,000 = 2.03 x 105 0.0009008 = 9.008 x 10 -4 0.00000001 = 6.34 x 10 -8 *c. “2” sextillion oxygen [O] atoms are in a drop of water and “4” sextillion hydrogen [H] atoms *1. (e.g.) Scientific notation: oxygen atoms = 2,000,000,000,000,000,000,000 = 2 x 1021 hydrogen atoms = 4,000,000,000,000,000,000,000 = 4 x 1021 *d. There are “1” quintillion ATOMS that make up one dot of ink on a newspaper *1. (e.g.) Scientific notation: 1,000,000,000,000,000,000 = 1 x 1018 *e. One sheet of paper is about 10,000 ATOMS in THICKNESS *1. (e.g.) Scientific notation: 10,000 = 1 x 104 f. A penny contains about “20” sextillion ATOMS of COPPER [Cu] and ZINC [Zn] 1. (e.g.) Scientific notation: 20,000,000,000,000,000,000,000 = 2 x 1022 B. Atomic Theory 1. Dalton’s THEORY of the ATOM having certain characteristics, form the BASIS of our understanding of ATOMS today a. All MATTER is composed of ATOMS b. Atoms can NOT be CREATED, DIVIDED or DESTROYED c. All the ATOMS of a specific ELEMENT are IDENTICAL 1. Explains why an ELEMENT always has the same PROPERTIES d. Atoms of two or more ELEMENTS can COMBINE to form NEW substances *e. Atoms of each ELEMENT have a unique MASS *1. The number of PROTONS plus NEUTRONS in the NUCLEUS give an atom its characteristic atomic MASS *f. The MASSES of the ELEMENTS in a COMPOUND are ALWAYS in the same RATIO *1. Since compounds have a specific chemical FORMULA, the ELEMENT(S) and the number (or RATIO) of ATOMS per ELEMENT are always the SAME. Therefore, the RATIO of the MASS of each ELEMENT is also always the SAME *a. (e.g.) calcium oxide [Ca5O2] = 5 (Ca amu) + 2 (O amu) = 232 amu 5 (40 amu) + 2 (16 amu) (200 amu) + (32 amu) 2. Thomson provided evidence of ATOMS having NEGATIVELY charged particles = ELECTRONS 3. Rutherford’s experiments provided evidence that the ATOM contained a DENSE, central NUCLEUS composed of POSITIVELY charged, SUBATOMIC particles called PROTONS 4. Chadwick discovered that the ATOM contained UNCHARGED particles called NEUTRONS 5. Bohr’s observations provided ATOM behavior based on ELECTRON movement around the NUCLEUS on circular ORBITAL paths (ENERGY levels) 6. Today: a. ELECTRON CLOUDS the area around the NUCLEUS where ELECTRONS move *b. QUARKS the sub-subatomic particles that make up PROTONS and NEUTRONS C. Parts of an Atom 1. Within the NUCLEUS: a. NUCLEUS: refers to the dense, inner most CORE of an ATOM containing the PROTONS (POSITIVELY charged particles) and the NEUTRONS (particles that have NO CHARGE) *1. the NUCLEUS is the most MASSIVE part of an ATOM b. PROTON: POSITIVE electrical CHARGE *1. written as: one proton = 1+; atom with “14” PROTONS = 14+ this MASSIVE subatomic particle along with the NEUTRON count gives an ATOM its atomic MASS number (measured in “amu” units or UNIFIED atomic mass units “u”) *1. mass of a PROTON = 0.0000000000000000000000017 u = 17 ten-septillionths u = 1.7 x 10-24 u (SCIENTIFIC NOTATION) number of PROTONS (atomic number) determines the IDENTITY of the ELEMENT *1. ATOMIC number is the PROTON count; unique for each ELEMENT *2. (e.g.) atomic number = 5; PROTON count = 5; element = BORON [B] proton COUNT can NOT change for an ELEMENT or you have a different ELEMENT *1. (e.g.) PROTON count = 10; element = NEON [Ne] “1” less proton = 9; element = FLOURINE [F] “1” more proton = 11; element = SODIUM [Na] c. NEUTRON: NO electrical CHARGE *1. (e.g.) beryllium [Be]: 4 protons = “4+” and 5 NEUTRONS = “4+” NET charge this MASSIVE subatomic particle along with the PROTON count gives an ATOM its atomic MASS number (measured in “amu” units or UNIFIED atomic mass “u”) *1. mass of a NEUTRON = slightly more than a PROTON *2. (e.g.) beryllium [Be]: 4 protons = “4+” and 5 NEUTRONS = “9” u (or “9” amu) neutron COUNT CAN change for an ELEMENT forming ISOTOPES *1. Steps to finding the NEUTRON count from a PERIODIC table *a. Find the atomic number (PROTON count) *b. Find the average atomic mass number (mixed decimal on Periodic Table) *c. ROUND the average atomic mass number to the nearest whole number *d. SUBTRACT the atomic number from the average atomic mass number *e. (e.g.) manganese [Mn] atomic number = 25 average atomic mass number = 54.94 ~ 55 manganese [Mn] = “30” neutrons *2. ISOTOPE atoms of the same ELEMENT with the same PROTON count (atomic number), but have a DIFFERENT number of NEUTRONS and therefore have a DIFFERENT atomic MASS number *a. hydrogen (protium) [H] proton = 1+; neutron = 0; 1 amu *b. hydrogen (deuterium) [H] proton = 1+; neutron = 1; 2 amu (isotope) *c. hydrogen (tritium) [H] proton = 1+; neutron = 2; 3 amu (isotope) 2. Outside the NUCLEUS: a. ELECTRON: NEGATIVE electrical CHARGE *1. written as: one electron = 1-; atom with “14” ELECTRONS = “14-” 2. the charges of PROTONS and ELECTRONS are OPPOSITE but EQUAL, which makes the ATOM electrically NEUTRAL (NET charge = “0”) *3. (e.g.) beryllium [Be]: 4 protons = “4+” + “4-” ELECTRONS = “0” net charge has relatively NO MASS; the ELECTRONS are NEVER used to determine the atomic MASS of an ELEMENT *1. MASS of an ELECTRON is 1,860 times LESS than a PROTON or NEUTRON *2. (e.g.) lithium [Li]: protons = 3+; electrons = 3-; neutrons = 4 = 7 u (or 7 amu) are in constant MOTION around the OUTSIDE of the nucleus within the electron CLOUD *1. HEISENBERG RULE states that it is NOT possible to determine the exact LOCATION and SPEED of an ELECTRON simultaneously *2. ELECTRON models show their movement on paths called energy ORBITALS (or energy LEVELS or energy SHELLS) electron COUNT CAN change for an ELEMENT when forming IONIC bonds by transferring (giving up or taking on) ELECTRONS, producing IONS *1. ION an electrically charged ATOM (having a “+” or “-” charge) due to ELECTRONS being TRANSFERRED (gained or lost) *2. IONIC bond the FORCE that attracts OPPOSITELY charged IONS and CHEMICALLY holds them together *a. sodium [Na] will lose “1” ELECTRON to become stable = “Na+” *b. chlorine [Cl] will gain “1” ELECTRON to become stable = “Cl-” *3. STABLE the outer-most, VALENCE energy ORBITAL contains the MAXIMUM number of valence ELECTRONS *a. helium [He] = 1st energy ORBITAL contains “2” electrons (maximum) *b. krypton [Kr] = 4th energy ORBITAL contains “8” electrons (maximum) D. Atomic Number verses Atomic Mass Number 1. Different COMBINATIONS of PROTONS, NEUTRONS and ELECTRONS produce ATOMS with different PROPERTIES (ELEMENTS) a. These different ATOMS chemically COMBINE to form different, NEW substances (COMPOUNDS) 2. The number of PROTONS distinguishes each ATOM from the other 118 ELEMENTS 3. ATOMIC NUMBER the number of PROTONS in the NUCLEUS of an ATOM a. (e.g.) iron [Fe] atomic number = 26; proton count = 26+ tin [Sn] atomic number = 50; proton count = 50+ mercury [Hg] (a LIQUID metal) atomic number = 80; proton count = 80+ bromine [Br] (only liquid NON-METAL) atomic number = 35; proton count = 35+ californium [Cf] (period 7; actinides) atomic number = 98; proton count = 98+ 4. The ATOM of a specific ELEMENT always has the SAME number of PROTONS, but it does NOT always have the same number of NEUTRONS a. (e.g.) chlorine [Cl] ALWAYS has 17+ PROTONS, but it could have 18 or 20 NEUTRONS b. chlorine-35 [Cl] with 17+ PROTONS and 18 NEUTRONS = ISOTOPE of CHLORINE chlorine-37 [Cl] with 17+ PROTONS and 20 NEUTRONS = ISOTOPE of CHLORINE 5. MASS NUMBER is TOTAL number of PROTONS plus NEUTRONS in the NUCLEUS of a specific ATOM a. (e.g.) silver [Ag] (period 5; group 11) proton count = 47+; neutron count = 61 = 108 u argon [Ar] (period 3; group 18) proton count = 18+; neutron count = 22 = 40 u arsenic [As] (period 4; group 15) proton count = 33+; neutron count = 42 = 75 u lead [Pb] (period 6; group 14) proton count = 82+; neutron count = 125 = 207 u uranium [U] (period 7; actinides) proton count = 92+; neutron count = 146 = 238 u *6. AVERAGE ATOMIC MASS is the weighted average of the MASSES of all the naturally occurring ISOTOPES of an ELEMENT *a. generally MIXED decimals on the Periodic Table *1. Exception: RADIOACTIVE elements’ most common ISOTOPES (whole numbers) b. (e.g.) potassium [K] (period 4; group 1) average atomic mass = 39.10 phosphorus [P] (period 3; group 15) average atomic mass = 30.97 gold [Au] (period 6; group 11) average atomic mass = 196.97 praseodymium [Pr] (period 6; lanthanides) average atomic mass = 140.91 mendelevium [Md] (period 7; actinides) average atomic mass = (258) II. LESSON 2 A. The History of the Periodic Table *1. In 1828, Dobereiner made one of the earliest attempts to “list” the ELEMENTS; proposed the Law of Triads stated there were groups of “3” ELEMENTS where the middle ELEMENT’S atomic MASS was the average of the other “2” ELEMENTS *a. (e.g.) calcium [Ca] = 40.08; strontium [Sr] = 87.62; barium [Ba] = 137.33 *2. In 1856, Newlands proposed the Law of Octaves stated that some of the 56 ELEMENTS whose atomic MASS differed by some multiple of EIGHT had similar PROPERTIES *a. this law supported the “Octet rule” for valence ELECTRONS and chemical BONDING 3. The first PERIODIC TABLE was created and published by MENDELEEV in 1869, who arranged the 63 existing ELEMENTS in order by their increasing ATOMIC MASS number a. “PERIODIC” means having a regular, REPEATING pattern and it means a “listing” b. Mendeleev used the “periodic patterns” to PREDICT future ELEMENTS 4. In 1914, Moseley re-organized the PERIODIC Table according to each element’s increasing ATOMIC number (PROTON count) from left to RIGHT rows (PERIODS), rather than by their atomic mass number 5. Today the PERIODIC TABLE is a valuable tool showing many PATTERNS among the 118 elements’ PROPERTIES *a. Some ELEMENTS are very reactive and form COMPOUNDS easily, while others are less REACTIVE, and still others do NOT form COMPOUNDS at all B. Information on the Periodic Table 1. Each SQUARE on the PERIODIC TABLE represents an ELEMENT 2. Each SQUARE generally includes the element’s: ATOMIC number corresponds to the number of PROTONS for that element *a. all ISOTOPES (different NEUTRON count) of that ELEMENT, but have the SAME ATOMIC number *b. all ELEMENTS on the periodic table are NEUTRAL (same number of PROTONS and ELECTRONS), the ATOMIC number also gives the ELECTRON count CHEMICAL symbol an ABBREVIATED form of the element’s NAME *a. 1st letter is always CAPITALIZED any other letter(s) are ALWAYS lower-case *b. some CHEMICAL symbols are taken directly from the ELEMENT name, others are based on their their LATIN, GREEK or Arabic names *c. some CHEMICAL symbols have “3” letters to represent their TEMPORARY name *1. (e.g.) yttrium [Y]; francium [Fr]; tungsten [W] (wolfram); Ununtrium [Uut] CHEMICAL name the name of the ELEMENT *a. sources for names come from SCIENTISTS; LOCATIONS; UNIVERSITIES; etc… *1. (e.g.) einsteinium [Es]; europium [Eu]; berkelium [Bk] AVERAGE ATOMIC MASS the weighted AVERAGE of the MASSES of all the natural occurring ISOTOPES of that ELEMENT *a. the unit label for ATOMIC MASS is “u” or “amu” *b. by rounding the AVERAGE ATOMIC MASS and subtracting the ATOMIC NUMBER from it, the NEUTRON count can be calculated for an ELEMENT *c. The ATOMIC MASS number of an ELEMENT can vary, because the ATOM of an ELEMENT can have varying numbers of NEUTRONS (ISOTOPES) *1. (e.g.) Carbon ATOMS must always contain 6 PROTONS, but they may contain anywhere from 5-8 NEUTRONS *d. ISOTOPES are ATOMS with the SAME number of PROTONS, but with DIFFERENT numbers of NEUTRONS. *1. (e.g.) Carbon-12 6 PROTONS and 6 NEUTRONS *a. Format: Atomic Mass Number 12 u Atomic Number 6 Carbon-14 6 protons and 8 NEUTRONS *a. Format: Atomic Mass Number 14 u Atomic Number 6 3. Inside the Square a. Xenon [Xe] square on the PERIODIC TABLE: 54 Xe xenon 131.29 ATOMIC NUMBER CHEMICAL SYMBOL CHEMICAL NAME AVERAGE ATOMIC MASS *C. Organization of the Periodic Table *1. The ELEMENTS are arranged in order of INCREASING ATOMIC numbers from LEFT to RIGHT and from TOP to BOTTOM *2. The ELEMENTS are arranged in “7” horizontal rows called PERIODS and “18” vertical columns called GROUPS (or FAMILIES) *a. Extensions of PERIOD 6 (lanthanides) & PERIOD 7 (actinides) are below the periodic table *3. An element’s PROPERTIES can be predicted based on its LOCATION on the periodic table *a. to the right, left or bordering the ZIGZAG line indicates the “classification” of the ELEMENT as a METAL, NON-METAL or METALLOID *b. the element’s PERIOD number and GROUP number indicates what “properties” the ELEMENT will have *4. The PERIODIC TABLE is “coded” (colors and symbols) to indicate the ELEMENTS’ STATE of matter (solid, liquid or gas) at ROOM temperature (250 C) and to assist in identifying its CLASSIFICATION (metal, non-metal or METALLOID) D. Metals, Non-metals and Metalloids 1. The “3” major CLASSIFICATION categories on the PERIODIC TABLE 2. The bolded, ZIGZAG line assists in the IDENTIFICATION as to what class the ELEMENT belongs a. METALS are to the LEFT of the LINE (Exception: hydrogen [H]) 1. PROPERTIES: LUSTER (shiny); CONDUCT electricity and HEAT; most are SOLID at 250 C; MALLEABLE (easily formed into different shapes) and DUCTILE (able to be made into a WIRE) b. NON-METALS are to the RIGHT of the ZIGZAG line 1. PROPERTIES: OPPOSITE of METALS (dull, POOR conductors of HEAT and electricity, BRITTLE) c. METALLOIDS border the ZIGZAG line 1. SIX metalloids: boron [B]; silicon [Si]; germanium [Ge]; arsenic [As]; antimony [Sb] and tellurium [Te] 1. PROPERTIES: have properties of both METALS and NON-METALS *a. (e.g.) Conduct HEAT: NOT as good as METALS; better than NON-METALS E. The Groups 1. GROUPS the “18” vertical COLUMNS forming the main body of the PERIODIC TABLE 2. The GROUPS have a FAMILY name, some based on the first ELEMENT in that COLUMN *a. Group 1 = ALKALI metals; Group 2 = ALKALINE EARTH metals; Groups 3-12 = TRANSITION metals (Group 11 = COINAGE metals); Group 13 = BORON family; Group 14 = CARBON family; Group 15 = NITROGEN family; Group 16 = OXYGEN family; Group 17 = HALOGENS; Group 18 = NOBLE gases (or inert gases) 3. The ELEMENTS in each GROUP have similar PHYSICAL and CHEMICAL properties because they have the SAME number of VALENCE electrons *a. VALENCE ELECTRON(S) are the ELECTRONS found on the OUTER-MOST portion of the electron CLOUD (energy orbital FARTHEST from the NUCLEUS) of an ATOM, which allows them to participate in chemical BONDING since they are farthest away from the attractive FORCE of the NUCLEUS *1. Group 1 are all METALS that react VIOLENTLY with WATER *a. From (Li – Fr) the alkali metals are ONLY found in COMPOUNDS *b. Group “1” valence electron count = “1” *c. Hydrogen [H] – which is NOT one of the alkali metals due to its different chemical PROPERTIES, is the simplest ELEMENT *1. 90% of the ATOMS of the universe are [H], but only 1% of the MASS of the Earth’s crust, OCEANS, and ATMOSPHERE is made up of [H] *2. Group 2, the alkaline earth metals, are good CONDUCTORS of electricity *a. From (Be – Ra) the alkaline earth metals = NOT as REACTIVE as Group 1 *b. Group “2” valence electron count = “2” *3. Groups 3-12, the transition metals, bridge the REACTIVE metals with the LESS reactive metals; good conductors of HEAT and ELECTRICITY; contain the most FAMILIAR metals; have LUSTER (shiny); and contains the iron [Fe] an essential METAL for the human BODY to make hemoglobin, which is necessary for carrying OXYGEN [O] in the bloodstream *a. transition metals do NOT follow the “Octet Rule” of needing “8” VALENCE electrons to be STABLE *b. STABLE an ATOM with the MAXIMUM number of ELECTRONS in its VALENCE orbital *c. Group 11, also known as COINAGE metals (copper [Cu], silver [Ag], gold [Au]), are slow to REACT with water *4. Groups 13-16 include METALS, NON-METALS, and the METALLOIDS and are composed of the BORON family (Group 13 w/ “3” valence electrons); CARBON family (Group 14 w/ “4” valence electrons); the NITROGEN family (Group 15 w/ “5” valence electrons) and the OXYGEN family (Group 16 w/ “6” valence electrons) *5. Group 17, the halogen family, reacts VIOLENTLY with the elements in GROUP 1, and therefore in their pure form are DANGEROUS to HUMANS, while in compounds they are very useful [(e.g.) Individually “Na” (explosive metal) and “Cl” (poisonous gas), but as a COMPOUND it forms TABLE SALT [NaCl] *a. Group “17” valence electron count = “7” *6. Group 18, the noble or inert gases, rarely REACT at all, since they are very stable ELEMENTS due to their VALENCE electron orbitals being FULL *a. Group “18” valence electron count = “8”; Exception: helium [He] = “2” F. The Periods 1. PERIOD is one of the 7 horizontal ROWS going across the PERIODIC TABLE containing a series of DIFFERENT types of ELEMENTS from different GROUPS (or families) a. The ELEMENTS in each PERIOD gradually change their physical and CHEMICAL PROPERTIES as you move from LEFT to RIGHT in PREDICTABLE ways: 1. (e.g.) atomic SIZE decreases (moving LEFT to RIGHT) 2. (e.g.) DENSITY of ELEMENTS is the LEAST dense on the LEFT and RIGHT sides of a PERIOD, while being MOST dense in the MIDDLE of the PERIOD *a. osmium [Os] (period 6; group8) has the HIGHEST known DENSITY *b. Period 1 contains two ELEMENTS: [H] and [He] and is referred to as a “short period” Periods 2 & 3 each contain 8 elements and are also “short periods” Periods 4 & 5 each contain 18 elements. Periods 6 & 7 each contain 32 ELEMENTS and therefore both have a portion which is SEPARATED from the main body of the PERIODIC TABLE *1. Period 6’s ELEMENTS (58-71) are the LANTHANIDES, which are MALLEABLE (easily shaped) and used to make ALLOYS (mixture of metals; metals & non-metals) *2. Period 7’s ELEMENTS (90-103) are the ACTINIDES, many are ARTIFICALLY made in labs, are unstable so they only last for a fraction of a SECOND after being made, and exist in small amounts (Exception: thorium (Th) and uranium [U]) *3. The LANTHANIDES and ACTINIDES are also known as the “rare earth elements” 4. The LANTHANIDES and ACTINIDES allow for the MAIN BODY of the PERIODIC TABLE to be NARROWER 5. The LANTHANIDES and ACTINIDES elements also are arranged in INCREASING order by their ATOMIC number *c. PERIOD number also indicates the number of energy ORBITALS an ATOM of a specific ELEMENT will have and for ELECTRON configuration its PRINCIPAL QUANTUM number Period 1 = hydrogen [H] and helium [He] only have “1” energy ORBITAL (type “s”) and their Principal QUANTUM number = “1” Period 2 has 2 energy ORBITALS (types: “s” and “p”); Principal quantum # = “2” Period 3 has 3 energy ORBITALS (types: “s” and “p”); Principal quantum # = “3” Period 4 has 4 energy ORBITALS (types: s; p and “d”); Principal quantum # = “4” Period 5 has 5 energy ORBITALS (types: s; p and “d”); Principal quantum # = “5” Period 6 has 6 energy ORBITALS (types: s; p; d and “f”); Principal quantum # = “6” Period 7 has 7 energy ORBITALS (types: s; p; d and “f”); Principal quantum # = “7” III. LESSON 3 A. Chemical Bonding and Chemical Changes 1. There are “118” elements and about “92” of these ELEMENTS are known to exist in NATURE and form all MATTER due to CHEMICAL BONDING 2. CHEMICAL BOND is a FORCE (or interaction) that holds ATOMS or IONS together forming MOLECULES 3. MOLECULE is the SMALLEST unit of a COMPOUND where two of more ATOMS are CHEMICALLY joined together by chemical BONDS a. some MOLECULES are extremely SMALL [(e.g.) water molecule [H2O] = “3” ATOMS] b. some are LARGE [(e.g.) DNA molecule [deoxyribonucleic acid] = “BILLIONS” of atoms] c. Due to the chemical BONDING of joining ATOMS together into a MOLECULE, the molecule acts as a single UNIT 4. CHEMICAL BONDING always involves chemical CHANGES by REARRANGING the ORDER of the ATOMS either by JOINING the atoms together or by BREAKING them apart *a. (e.g.) 2 H2 + O2 2 H2O (SYNTHESIS chemical reaction) “2” diatomic MOLECULES of HYDROGEN gas COMBINE with “1” DIATOMIC MOLECULE of OXYGEN gas, which YIELDS “2” MOLECULES of WATER *b. (e.g.) 2 H2O 2 H2 + O2 (DECOMPOSITION chemical reaction) “2” MOLECULES of WATER (decompose), which YIELDS “2” DIATOMIC MOLECULES of HYDROGEN gas AND “1” diatomic MOLECULE of oxygen GAS *5. When substances undergo a CHEMICAL change, NEW substances with different PHYSICAL and chemical PROPERTIES are formed *a. (e.g.) The properties of the REACTANTS are completely DIFFERENT from those of the PRODUCT, yet hydrogen [H] and oxygen [O] are the only ELEMENTS used in the formation of the 2 water MOLECULES 2 H2 + O2 (the REACTANTS) 2 H2O (the product) *b. The CHEMICAL change did NOT create or DESTROY any ATOMS (Law of CONSERVATION of MATTER), it simply caused the hydrogen [H] and oxygen [O] ATOMS to be REARRANGED Balancing the equation: 2 H2 + O2 (the REACTANTS) 2 H2O (the PRODUCT) Reactants (original substances): Elements: Number of atoms/element: H = 4 O = 2 Elements: H O Products (new substances): Number of atoms/element: = 4 = 2 B. Atomic Models *1. Using Models *a. MODEL is a representation of how something LOOKS and/or WORKS *b. The PARTICLE model of MATTER is used to study ATOMS and MOLECULES *1. The study of ATOMS explains the various PROPERTIES of ELEMENTS *2. The study of ELEMENTS explains the various PROPERTIES of COMPOUNDS 2. MODELS of ATOMS are used to show how atoms BEHAVE and how ELECTRONS are involved in chemical BONDING 3. There are “3” different MODELS used to represent ATOMS: ELECTRON CLOUD model, the BOHR model and the SPACE-FILLED model a. ELECTRON CLOUD model shows the atomic NUCLEUS in the CENTER of the atom and the indistinct, CLOUD-like region around the NUCLEUS 1. ELECTRONS move through this CLOUD-like region *2. Model re-enforces the HEISENBERG rule that the POSITION and trajectory of moving ELECTRONS can NOT be precisely determined at the SAME instant 3. ELECTRON CLOUD models do NOT show the NUMBER of ELECTRONS b. BOHR model shows the number of ENERGY levels and the number of ELECTRONS per ENERGY level (or energy ORBITAL) 1. ELECTRONS are represented by DOTS 2. ENERGY levels are represented by RINGS 3. The BOHR model does NOT show the TRUE location of the ELECTRONS, but it helps explain the chemical BONDING of ATOMS and their chemical PROPERTIES c. SPACE-FILLED model represent ATOMS as SOLID SPHERES to show how they are JOINED together in substances 1. The SPACE-FILLED model does NOT show the PARTS that make up the ATOMS C. Valence Electrons and Chemical Bonds *1. VALENCE electrons explain BONDING power because they can either be SHARED or TRANSFERRED (lost/GAINED) between other ATOMS *2. BONDING power refers to the number of chemical BONDS an ELEMENT can form during a CHEMICAL change 3. VALENCE ELECTRON are the ELECTRONS that occupy the atom’s OUTER-MOST energy level (valence electron ORBITAL) and have the LEAST amount of ATTRACTIVE force TO the NUCLEUS *a. They are the ONLY electrons that can be SHARED or TRANSFERRED (lost/GAINED) *b. The AMOUNT of VALENCE electrons in the VALENCE electron orbital determines whether the ATOM of an ELEMENT will give UP, take ON or SHARE electrons *c. The number of VALENCE electrons an element has INCREASES from LEFT to RIGHT across a PERIOD 4. VALENCE ELECTRON ORBITAL is the atom’s OUTER-MOST energy orbital that is FARTHEST from the ATTRACTIVE force TO the NUCLEUS *5. Guidelines to follow for the amount of ELECTRONS placed on energy ORBITALS are: *a. 1st energy orbital = “2” electrons (maximum); 2nd energy orbital = “8” electrons (maximum); 3rd energy orbital = “18” electrons (maximum); 4th energy orbital = “32” electrons (maximum); 5th energy orbital = “64” electrons (maximum); etc… *b. The “18-Electron Rule”: 1st energy orbital = “2” electrons (max); 2nd energy orbital = “8” electrons (maximum); 3rd energy orbital = “18” electrons (maximum); all other energy orbitals “max” out at “18” electrons and are then considered STABLE *c. The “Octet Rule” (or LEWIS RULE of Eight): states that ATOMS combine to form MOLECULES by losing, GAINING or sharing VALENCE electrons until they attain “8” VALENCE electrons and are considered STABLE: 1st energy orbital = “2” electrons (maximum); 2nd energy orbital = “8” electrons (maximum); all other energy orbitals “max” out at “8” electrons 6. The PERIODIC table (using the OCTET Rule) can also indicate the number of VALENCE electrons for an ELEMENT a. The ELEMENTS in each GROUP of the periodic table have the SAME NUMBER and arrangement of VALENCE ELECTRONS in their VALENCE electron orbital b. The GROUP number in GROUPS 1-2 and 13-18 can be used to find the NUMBER of VALENCE electrons in their VALENCE electron orbital 1. EXCEPTIONS: a. the transition metals (GROUPS 3-12) have either 1 or 2 VALENCE electrons b. helium [He] is in GROUP 18, but only has “2” VALENCE electrons due to its single VALENCE ELECTRON ORBITAL 2. Group 1 (ALKALI metals) - all the ELEMENTS have “1” valence electron [H = SHARES; Li-Fr = TRANSFER by GIVING it up] Group 2 (alkaline earth metals) - all the ELEMENTS have “2” valence electrons [TRANSFER by GIVING them up] Groups 3-12 (transition metals) - most the ELEMENTS have 1 or 2 valence electrons [TRANSFER by GIVING them up] Group 13 (BORON family) - all the ELEMENTS have “3” valence electrons [TRANSFER by GIVING them up] Group 14 (CARBON family) - all the ELEMENTS have “4” valence electrons [SHARE; especially CARBON] Group 15 (NITROGEN family) - all the ELEMENTS have “5” valence electrons [TRANSFER by TAKING on electrons given up by other atoms] Group 16 (OXYGEN family) - all the ELEMENTS have “6” valence electrons [TRANSFER by TAKING on electrons given up by other atoms] Group 17 (HALOGEN family) - all the ELEMENTS have “7” valence electrons [TRANSFER by TAKING on electrons given up by other atoms] Group 18 (NOBLE or INERT gases) - all the ELEMENTS have 8 valence electrons (EXCEPTION: HELIUM [He] has “2” valence electrons) [All STABLE elements; do NOT TRANSFER or SHARE their valence electrons] 7. ATOMS want to be STABLE (valence orbital is FULL) but if their VALENCE orbital does NOT contain the maximum number of VALENCE electrons they will form chemical BONDS with other ATOMS by TRANSFERRING (gain/lose) or SHARING VALENCE electrons a. Forming BONDS allows ATOMS to FILL their VALENCE orbitals *b. (e.g.) Group 1 (ALKALI metals) having “1” VALENCE electrons (very REACTIVE elements) want to BOND with elements from Group 17 (HALOGEN family) that have “7” VALENCE electrons (LESS reactive elements) c. (e.g.) sodium [Na] (Group 1) has “1” VALENCE electron (LOSE “1” electron) = Na+ chlorine [Cl] (Group 17) has “7” VALENCE electrons (GAIN “1” electron) = ClCHEMICAL bond (ionic bond) by TRANSFERRING = NaCl (IONIC compound) *D. Valence Electrons and Electron Configuration *1. ELECTRON CONFIGURATION a term for how ELECTRONS are arranged in an ATOM *2. ELECTRON movement around the NUCLEUS is NOT in circular paths (Bohr’s model), but in CLOUD-like zones (ORBITALS) *3. HEISENBERG Uncertainty Rule states that it is IMPOSSIBLE to know the LOCATION and SPEED of the ELECTRONS simultaneously *4. “ORBITALS” are where the ELECTRONS are LIKELY to be FOUND *a. ORBITALS are GROUPED and IDENTIFIED according to their SHAPES: s, p, d and f (also known as SUBSHELLS) “s” “sharp” “p” “principal” “d” “diffuse” “f” “fundamental” *5. QUANTUM numbers the “4” numbers that describe the atomic ORBITALS: n, l, ml, ms *a. PRINCIPAL QUANTUM number (n) is the WHOLE number which tells the overall ENERGY of the ELECTRON and its LOCATION “n” is located in front of the PERIOD number and is the SAME as the PERIOD # the HIGHER the PRINCIPAL QUANTUM number the MORE electron energy FARTHER away from the NUCLEUS the MORE electron energy ELECTRONS with more ENERGY can occupy different types of ORBITALS LOWER energy ORBITALS fill FIRST (Exception: TRANSITION metals) *b. QUANTUM number (l) represents the orbital SHAPE and refers the “GROUPS” (1-18) on the PERIODIC table Type “s” = spherical shape; “s” block = GROUPS “1 & 2” and HELIUM [He] Type “p” = dumbbell shape; “p” block = GROUPS “13 - 18” (without [He]) Type “d” = many complicated shapes; “d” block = GROUPS “3 - 12” Type “f” = more complicated shapes; “f” block = (lanthanides & actinides) *c. QUANTUM number (ml) represents the orientation of the ORBITAL which only effects the “p, d and f” ORBITALS *d. QUANTUM number (ms) represents the SPIN of the ELECTRON on its axis in one of two possible directions *1. An ORBITAL can only accommodate “2” ELECTRONS each since they are SPINNING in an OPPOSITE direction *2. Each ORBITAL can HOLD a MAXIMUM (based on “2” ELECTRONS per ORBITAL ms) of: Type “s” = has “1” ORBITALS = TOTAL “2” ELECTRONS Type “p” = has “3” ORBITALS = TOTAL “6” ELECTRONS Type “d” = has “5” ORBITALS = TOTAL “10” ELECTRONS Type “f” = has “7” ORBITALS = TOTAL “14” ELECTRONS Orbital Type [quantum # (l)] s Number of Orbitals [(m2) max # of 2 electrons per subshell] 1 Maximum # of Electrons per Type of Orbital 2 p 3 6 d 5 10 f 7 14 *3. HUND RULE of MAXIMUM MULTIPLICITY (ms) states that you MUST put “1” ELECTRON in each ORBITAL of a subshell BEFORE you begin to DOUBLE (or PAIR) the ELECTRONS up *a. Each “p” ORBITAL (x, y and z) MUST receive “1” ELECTRON each BEFORE electron “PAIRING” can occur *b. (e.g.) oxygen [O] ATOMIC number (PROTON count) = “8” ELECTRON count (NEUTRAL atom) = “8” Each ARROW represents “1” ELECTRON 1s 2s 2p (x) 2p (y) 2p (z) *6. Writing FORMAT of ELECTRON CONFIGURATION for an ELEMENT: n, l, ml, ms *a. 1st write the PRINCIPAL QUANTUM number (“n”) which represents the energy ORBITAL ( or “PERIOD” number) nd *b. 2 write the ORBITAL letter (“l”) representing the “ORBITAL type” (s, p, d or f) *c. 3rd write the “EXPONENT” which represents the number of “ELECTRONS” in that ORBITAL *d. The “ORDER” for filling ORBITALS: 1s, 2s, 2p, 3s, 3p, 4s… For the order of filling orbitals, begin at the base of each arrow starting with “1 s”, follow it all the way to the point and begin at the base of the next arrow. Order would be: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, etc… n1 n2 n3 n4 n5 n6 n7 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 7p 3d 4d 5d 6d 4f 5f *e. (e.g.) Writing FORMAT of ELECTRON CONFIGURATION for: *1. lithium [Li] ATOMIC number (PROTON count) = “3” ELECTRON count (NEUTRAL atom) = “3” ELECTRON CONFIGURATION: ___1s2 2s1_______________ (drawing) *2. carbon [C] ATOMIC number (PROTON count) = “6” ELECTRON count (NEUTRAL atom) = “6” ELECTRON CONFIGURATION: ____1s2 2s2_2p2___________ (drawing) *3. sodium [Na] ATOMIC number (PROTON count) = “11” ELECTRON count (NEUTRAL atom) = “11” ELECTRON CONFIGURATION: ____1s2 2s2_2p6_3s1_______ (drawing) *8. The VALENCE electrons in ELECTRON CONFIGURATION refer to ALL the electrons in the HIGHEST PRINCIPAL QUANTUM number ORBITAL(S) *a. (e.g.) nitrogen [N] ATOMIC number (PROTON count) = “7” ELECTRON count (NEUTRAL atom) = “7” Electron Configuration: 1s2 2s 2 2p 3 Valence electron count: “5” *b. (e.g.) sodium [Na] ATOMIC number (PROTON count) = “11” ELECTRON count (NEUTRAL atom) = “11” Electron Configuration: 1s2 2s2 2p6 3s 1 Valence electron count: “1” *c. For GROUPS 1, 2, 13-18 (EXCEPTION: helium [He]) the GROUP number will help to determine the VALENCE ELECTRON count *1. (e.g.) lithium [Li] = Group “1” “1” valence electron carbon [C] = Group “14” “4” valence electron neon [N] = Group “18” “8” valence electron magnesium [Mg] = Group “2” “2” valence electron helium [He] = Group “18” “2” valence electron (exception) *9. Information that can be determined by reading the ELECTRON CONFIGURATION of an ELEMENT: ATOMIC number (or PROTON count); ELECTRON count (NEUTRAL atom); PRINCIPAL QUANTUM (or PERIOD number); VALENCE electron count (or GROUP number); NEUTRAL (same number of PROTONS as ELECTRONS); STABLE ( VALENCE orbital is FULL); Type of BOND (IONIC, COVALENT or INERT) *a. (e.g.) sulfur [S] Electron Configuration: 1s2 2s2 2p6 3s2 3p4 Atomic number/Proton or Electron count (sum of ALL the exponents) = “16” Principal Quantum or Period number (highest LARGE number) = “3” Valence Electron count (sum of exponents of highest LARGE number) = “6” Group number (sum of exponents of highest LARGE number) = “6” Neutral (when proton and electron count are the SAME number) = “YES” Stable (when valence electron count is “8” [“2” for helium]) = “NO” Type of Bond (ionic: transfer; covalent: share; inert: Grp18) = “IONIC”