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The Atom Chapter 2 Chem2A 1 2 3 Elements • An element is a pure substance that cannot be broken down into simpler substances by a chemical reaction. • Consists of a dense core, surrounded by a negatively charged cloud • Contains three types of subatomic particles – Proton – Neutron – Electron 4 The Proton • Charge = +1.602x10-19 C – +1 • Located in the nucleus • Mass = 1.672 x 10-24 g – ~1 atomic mass unit (amu) – 1 amu = 1.66 x 10-24 g • 1/12 the weight of a 12C atom 5 The Neutron • No charge (O C) • Located in the nucleus • Mass = 1.675 x 10-24 g – ~1 amu 6 The Electron • Charge = -1.602 x 10-19 C – -1 • Located outside the nucleus in an e- “cloud” • Mass = 9.109 x 10-28 g – ~0 amu 7 Plum Pudding Model • J.J. Thomson • 1904 8 Gold Foil Experiment • Ernest Rutherford • 1909 9 Thomson vs. Rutherford • Rutherford used the angle of deflection to calculate size of the nucleus • Nucleus approximated at 10,000 times smaller than the radius of the atom • Atom is actually mostly empty space 10 Atoms 11 12 13 Periods 14 Groups 15 Periodic Table of Elements 16 Classification of Periodic Table 17 Elements and Symbols • • • • • C = Carbon N = Nitrogen O = Oxygen Cl = Chlorine Ba = Barium • • • • • U=? Cf = ? Bk = ? Md = ? Es = ? • • • • K= Pb = W= Sb = • • • • • Co = Cu = Cr = Mg = Mn = Metals • Shiny •Conduct Electricity •Ductile – Can be drawn into a thin wire •Malleable (Shapeable) •High M.P. and B.P. •Solids @ RT – Except Hg 20 Non-Metals • • • • • Don’t conduct well Not usually ductile Brittle Low M.P. and B.P. Many are gasses at RT 21 Metalloids • Chemical characteristics in between metals and non-metals • Includes all atoms falling on the dividing line between metals and non metals 22 Classification of Periodic Table 23 Alkali Metals • • • • • Group 1 (1A) Li, Na, K etc. Soft, shiny metals Conduct heat and electricity React violently with H2O – Form H2(g) and alkaline (basic) solutions Akali(ne) Earth Metals • • • • Group 2 (2A) Be, Mg, Ca etc. Not as reactive as Alkali Metals, but still quite reactive Tend to make basic solutions when placed in water Transition Metals • Groups 3-12 • Tend to have high densities and B.P. • All are metals • Often used for electrical conduction • Often have vivid colors when in solution – Used for pigments Colors of Transition Metal Compounds Iron Cobalt Nickel Copper Zinc Lanthanides • Elements 57-71 – Lanthanum (La) to Lutetium (Lu) • Commonly used in lasers • Can deflect UV and infrared rays Actinides/Actinoids • Elements 89-103 – Actinium (Ac) to Lawrencium (Lr) • Only Actinium, Thorium (Th), and Uranium (U) occur naturally – Others created by neutron bombardment • Radioactive Groups 13(3A) – 16(6A) • No common name • Boundary between metals and non-metals occurs here • Contain elements abundant in earth’s crust, atmosphere, and living things • Contains the metalloids Halogens • Group 17 (7A) • Love to form salts with metals – NaCl, KBr, CaCl2 • Like to form diatomic molecules – F2, Cl2, Br2 Noble Gases • Group 18 (8A) • Very unreactive • Don’t like to bond to other molecules • Generally not abundant Diatomic Molecules Dinitrogen (N2) • Molecules consisting of only two atoms of either the same or different elements – O2 – CO • Homonuclear Diatomic Molecule: a molecule made of two atoms of the same element – H2 • Heteronuclear Diatomic Molecule: a molecule made of two atoms that are different elements – NO Allotropism • The existence of multiple pure forms of an element, in the same phase (solid, liquid, or gas), that differ in structure • Different forms are called allotropes • Can exhibit varied physical properties and chemical behaviors • Don’t confuse allotropes with isotopes!!! O2 O3 2.2 Structure of the Atom From the periodic table: Atomic Symbol Atomic Number = # of protons Atomic Mass • Any given element is neutral • # of protons (+) = # of electrons (-) • # of neutrons = Atomic Mass (rounded) - # Protons – 7 – 3 = 4 neutrons 36 Problems H C Fe Pd W Atomic Weight # Protons # Electrons # Neutrons 37 Calculator Boot Camp 38 2.3 Isotopes A. Isotopes, Atomic Number, and Mass Number Isotopes are atoms of the same element that have a different number of neutrons. the number of protons (Z) + the number of neutrons Mass number (A) = Mass number (A) 35 Atomic number (Z) 17 # of protons Cl 37 17 Cl Mass number (A) Atomic number (Z) = 17 # of protons = 17 # of electrons = 17 # of electrons = 17 # of neutrons = 35 – 17 = 18 # of neutrons = 37 – 17 = 20 39 Identify the atomic mass, number of protons, and number of neutrons for the following elements. 2.3 Isotopes B. Atomic Weight The atomic weight is the weighted average of the masses of the naturally occurring isotopes of a particular element reported in atomic mass units. From the periodic table: 6 C 12.01 atomic number element symbol atomic weight (amu) 41 2.3 Isotopes B. Atomic Weight HOW TO Determine the Atomic Weight of an Element Example What is the average atomic weight of chlorine? Step [1] List each isotope, its mass in atomic mass units, and its abundance in nature. Isotope Mass (amu) Isotopic Abundance Cl-35 34.97 75.78% = 0.7578 Cl-37 36.97 24.22% = 0.2422 42 2.3 Isotopes B. Atomic Weight HOW TO Determine the Atomic Weight of an Element Step [2] Multiply the isotopic abundance by the mass of each isotope, and add up the products. The sum is the atomic weight of the element. 43 1. Using the percent abundances below, calculate the average atomic mass for Carbon a) C-12 = 98.890% b) C-13 = 1.110% 2. Antimony (Sb) has two stable isotopes, 121Sb and 123Sb with masses of 120.9038 u and 122.9042 u, respectively. Calculate the percent abundances of these two isotopes 45 Atomic Orbitals and Electron Configurations Electronic Structure •Electrons are confined to discrete regions called shells •The shells are numbered n = 1, 2, 3, etc. •Electrons in lower numbered shells are closer to the nucleus and are lower in energy. •Electrons in higher numbered shells are further from the nucleus and are higher in energy. 47 Electronic Structure • Shells with larger numbers (n) are farther from the nucleus and can hold more electrons. • # of electrons in each shell is calculated using 2(n2), where n = shell number. Shell (n) increasing energy Number of Electrons in a Shell 4 32 3 18 2 8 1 2 increasing number of electrons 48 Electronic Structure • Shells are divided into subshells, identified by the letters s, p, d, and f. • The subshells consist of orbitals. • An orbital is a region of space where the probability of finding an electron is high. • Each orbital can hold two electrons. increasing energy Subshell s p Number of Orbitals 1 3 d 5 f 7 49 S Orbitals • Spherical shape • Lowest energy of the orbitals • Contains 1 orbital: s – 1 orbital can hold 2 electrons MAX – = 2 electrons max 50 P Orbitals • The p orbital has a dumbbell shape • Higher energy than S Orbitals • 3 separate orbitals px, py, pz – Each orbitals can hold 2 electrons MAX – = 6 electrons max 51 D Orbitals • Double dumb bell shape or single dumb bell with a donut • Higher energy than s and p • 5 orbitals • dxy, dxz, dyz, dx2-y2, dz2 • = 10 electrons max 52 F Orbitals • • • • Complex Shapes Highest Energy Orbital 7 Orbitals =14 electrons max 53 54 Electronic Structure 55 Electron Configuration • The electron configuration shows how the electrons are arranged in an atom’s orbitals. • The ground state is the lowest energy arrangement. Rules to Determine the Ground State Electronic Configuration of an Atom Rule [1] •Electrons are placed in the lowest energy orbital beginning with the 1s orbital. •Orbitals are then filled in order of increasing energy. 56 Electron Configuration Rules to Determine the Ground State Electronic Configuration of an Atom 57 Electron Configuration Rules to Determine the Ground State Electronic Configuration of an Atom Rule [2] Each orbital holds a maximum of 2 electrons. Rule [3] When orbitals are equal in energy: •1 electron is added to each orbital until all of the orbitals are half-filled. •Then, the orbitals can be completely filled. 58 Electron Configuration Orbital Diagrams An orbital diagram uses a box to represent each orbital and arrows to represent electrons. an orbital a single, unpaired electron an electron pair Two electrons must have paired spins (opposite directions) to fit into the same orbital. 59 Electron Configuration A. First-Row Elements (Period 1) Element Orbital Notation Electron Configuration H (Z = 1) 1 electron He (Z = 2) 2 electrons 60 Electron Configuration B. Second-Row Elements (Period 2) Element Orbital Notation Electron Configuration Li (Z = 3) 3 electrons C (Z = 6) 6 electrons Ne (Z = 10) 10 electrons 61 Problems • Determine the electron configuration and orbital diagrams for the following atoms 1. N 2. O 3. Na 4. Mg 5. Fe 6. W 62 Electron Configuration • The electron configuration can be shortened by using Noble Gas Notation. • Write the Symbol of the previous Noble Gas, then add the electronic configuration of the additional electrons. element: C previous noble gas: 63 Electron Configuration C. Other Elements Element Ca 20 electrons 64 Electron Configurations and the Periodic Table 65 2.7 Valence Electrons • The chemical properties of an element depend on the number of electrons in the valence shell. • The valence shell is the outermost shell (the highest value of n). • The electrons in the valence shell are called valence electrons. Be Cl 1s22s2 1s22s22p63s23p5 valence shell: n = 2 valence shell: n = 3 # of valence electrons = 2 # of valence electrons = 7 66 2.7 Valence Electrons A. Relating Valence Electrons to Group Number • Elements in the same group have similar electron configurations. • Elements in the same group have the same number of valence electrons. • The group number, 1A–8A, equals the number of valence electrons for the main group elements. • The exception is He, which has only 2 valence electrons. • The chemical properties of a group are therefore very similar. 67 2.7 Valence Electrons A. Relating Valence Electrons to Group Number Group number: 1A 2A 3A Period 1: H 1s1 4A 5A 6A 7A 8A He 1s2 Period 2: Li 2s1 Be B C N O F Ne 2s2 2s22p1 2s22p2 2s22p3 2s22p4 2s22p5 2s22p6 Period 3: Na Mg Al Si P S Cl Ar 3s1 3s2 3s23p1 3s23p2 3s23p3 3s23p4 3s23p5 3s23p6 Valence Electrons B. Electron-Dot Symbols • Dots representing valence electrons are placed on the four sides of an element symbol. • Each dot represents one valence electron. • For 1 to 4 valence electrons, single dots are used. With more than 4 valence electrons, the dots are paired. Element: H C O Cl # of Valence electrons: 1 4 6 7 Electron-dot symbol: H C O Cl 69 Periodic Trends A. Atomic Size Increases •The size of atoms increases down a column Decreases •The size of atoms decreases across a row 70 Chapter 2 Review • The periodic table will be folded up for quiz and exam • Know the names and elemental symbols of s block, p block, and select d block atoms – W, Cu, Ag, Fe, Hg, Ti 71 Chapter 2 Review • Know how to classify atoms as metals, non-metals, and metalloids. • Know the period classifications – Alkali, Alkaline, Transition Metals, main block, halogens, noble gas 72 Classification of Periodic Table 13 Al 84 Po 73 Chapter 2 Review • Location, and charge of the 3 particles that make up an atom – Proton (nucleus), Neutron (nucleus), Electron (diffuse cloud surrounding nucleus) • Determining how many protons, neutrons and electrons an atom has 74 Chapter 2 Review • Orbitals – s, p, d, f – Know everything about s and p – Know how many electrons all can have – Know how they relate in energy 75 Chapter 2 Review • What are isotopes • How do isotopes relate in terms of weight • How to calculate the average atomic weight 76 Chapter 2 Review • Know how to fill in (yes ill give you boxes) molecular orbitals • Know how to write noble gas abbreviations for any atom 77 Chapter 2 Review • Know what valence electrons are and how to ID them in an electron configuration • Know the relationship between atoms in the same period • Know how to write electron dot symbols 78