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REVIEW 2 Atoms and Elements 4 4.1 ELEMENTS A. Atoms make up all matter B. ~91 different naturally occurring elements 1. Each element is composed of tiny.

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Presentation on theme: "REVIEW 2 Atoms and Elements 4 4.1 ELEMENTS A. Atoms make up all matter B. ~91 different naturally occurring elements 1. Each element is composed of tiny."— Presentation transcript:

1 REVIEW 2 Atoms and Elements ELEMENTS A. Atoms make up all matter B. ~91 different naturally occurring elements 1. Each element is composed of tiny indestructible particles called atoms 2. All atoms of a given element have the same mass and other properties that distinguish them from atoms of other elements 3. Atoms combine in simple, whole-number ratios to form compounds 4.2 Names vs symbols of elememts- Memorize List from web. Recognize Metal, non-metals, metalloids -

2 4.3 Dalton’s Atomic Theory A. Atoms make up all matter B. ~91 different naturally occurring elements Indivisble: The Atomic Theory A. Democritus (~400 BC) B. Dalton (early 1800's) 1. Each element is composed of tiny indestructible particles called atoms 2. All atoms of a given element have the same mass and other properties that distinguish them from atoms of other elements 3. Atoms combine in simple, whole-number ratios to form compounds -Law of conservation of matter 4.4 Formulas of Compounds- More detail in Chapter 5

3 The Atomic Structure and Nucleus A. Positive and negative charges (Thompson)-plum pudding model B. Mostly empty space (Rutherford) C. Current nuclear theory of the atom 1. Most of the atom's mass and all of its positive charge is contained in the nucleus 2. Most of the atom is empty space with small fast moving electrons 3. Number of protons = number of electrons in neutral atoms The Properties of Protons, Neutrons, and Electrons -Atomic Number = Number of protons - Mass Number = # of Protons + # of neutrons A. Masses expressed in atomic mass units (amu) B. Protons and neutrons about 2,000 more massive than electrons C. Nature of electrical charge 1. Electrical charge is a fundamental property of protons and electrons 2. Species of opposite electric charge are attracted to each other 3. Species of the same electric charge repel each other 4. Species of opposite electric charge cancel to make a neutral atom

4 4.7 Isotopes Isotopes: When the Number of Neutrons Varies in an element A. Isotopes have the same chemical properties, but different masses B. Some isotopes are more prevalent than others C. Chemical symbol with mass number indicates which isotope D. Some elements have many isotopes, some very few Elements: Defined by Their Number of Protons A. Identity of atom comes from the number of protons B. Atomic number (Z) = number of protons C. Chemical symbol is a shorthand notation of chemical name 1. C = carbon = 6 protons 2. O = oxygen = 8 Ptrotons 3. Co = cobalt = 27 Protons Atomic Mass A. Value on periodic table represents average mass of all isotopes B. No chlorine atom weighs exactly amu 35 Cl 17 Mass #At. #

5 4.8 Looking for Patterns: The Periodic Law and the Periodic Table A. Mendeleev ( ) B. Periodic law C. Metals (usually loose electrons to form cations) D. Nonmetals (usually gains electrons to form anions) E. Metalloids, also known as semiconductors F. Individual group names 1. Group 1 – alkali metals +1 ion only 2. Group 2 – alkali earth metals +2 ion only 3. Group 7 – halogens usually –1 ion 4. Group 8 – noble gases Elements which are diatomics N 2 O 2 F 2 Cl 2 Br 2 I 2 and H Ionic compound – It must contain a metal, Meatl forms + ion CATIONS Nonmatels form – charge ions ANIONS Number of atoms in a chemical formula

6 Chapter 11- Modern Atomic Theory 11.2 Light: Electromagnetic Radiation A. Light is not matter, it is pure energy B. A photon is a particle of light 1. Wavelength is distance between adjacent wave crests 2. Frequency is the number of cycles that pass a point per unit time The Electromagnetic Spectrum: Memorize the order (starting in low energy from Radio waves, Microwaves, Infrared, visible, UV, X-rays, Gamma rays) A. Gamma rays 1. High energy 2. Low wavelength B. X-rays C. Ultraviolet D. Visible 1. Very narrow band 2. Does not affect biological molecules E. Infrared F. Microwaves G. Radio waves 4.2 Classical vs Quantum physics for matter and light

7 The Bohr Model: Atoms with Orbits A. Electrons exist in quantized Circular orbits B. Electrons get excited to a higher energy level by absorbing energy C. Electrons emit light when relaxing from higher levels D. Absorbed and emitted light correspond to energy difference between states E. Higher the principal shell# higher is the energy of the electron (away from nucleus) The Quantum-Mechanical Model: Atoms with Orbitals A. Orbitals instead of Circular orbits B. Electrons do not move in predictable paths C. Probability maps-different shapes- called as ORBITALS Quantum-Mechanical Orbits (shapes) A. Quantum numbers 1. Principal shell 2. Subshell: s (2e), p(6e), d(10e), f(14e) B. Electronic configuration 1. Gives location of each electron in the atom 2. Li: 1s 2 2s 1 C. How do the electrons fill in the atom 1. Auf-Bau Principle, Hund’s rule 2. Electron spin-Pauli exclusion principle Filling order 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s-MEMORIZE

8 11.10 Electron Configurations and the Periodic Table A. Valence electrons (outermost shell electrons with highest principal shell#) B. Core electrons C. Writing electronic configurations 1. Noble gas core 2. Locations of valence electrons 3. Highest n value is equal to row value 4. d electron group is equal to row number – Periodic Trends: Ionization Energy, Atomic Size, and EN A. Atomic size 1. Decreases as you move to the right along a row (period) 2. Increases as you go down a column (group) Ionization energy has trends opposite to Atomic Size Compound Formation A. Valence electrons dictate chemical reactivity B. Elements in the same group have similar reactivity as they have SAME VALENCE ELCTRONIC CONFIGURATION C. Compound formation and Octet rule with Chapter 5

9 Chapter Electronegativity and Polarity: A. Electronegativity B. Bond dipoles C. Pure covalent bonds D. Polar and nonpolar bonds E. Ionic bonds 12.5 Ionic compounds A. Ionic compounds formed between a metal and one or more nonmetals Writing Formulas for Ionic Compounds A. Ionic compounds always contain positive and negative ions B. Formula units form neutral specie Compounds held by covalent bonds Covalent Bond -Molecular compounds formed between two or more nonmetals - Attractive forces maximum at bond position - Sharing of electrons A. Every sample of a compound has the same proportions of constituent elements A Molecular View of Elements and Compounds A. Atomic elements exist in nature as single atoms B. Molecular elements exist as diatomic molecules, H 2, N 2, O 2, F 2, Cl 2, Br 2, I 2

10 Representing Valence Electrons with Dots () A. Valence electrons in outermost shell electrons B. Lewis Structure is simply chemical symbol + # valence electrons around it C. Most atoms "want" eight electrons around them (octet rule) D. Exception: duet rule for hydrogen and helium E. A covalent bond is a shared pair of electrons to get an octet around each atom 12.8 Covalent Lewis Structures: Electrons Shared A. Two bonded nonmetals share electrons such that both get an octet (duet) B. Both species get "credit" for all electrons in bond C. Two species can share two, four, or six electrons 1. Two electrons shared is a single bond 2. Four electrons shared is a double bond 3. Six electrons shared is a triple bond

11 12.8 Writing Lewis Structures for Covalent Compounds A. Identify central atom, write skeletal structure for the molecule (H,F never in center) B. Calculate the total number of valence electrons, C. make single bonds, take away electrons for single bond C. Distribute remaining electrons to corner atoms giving each an octet (duet) D. Further remaining electron give it to center D. If center still the lacks octet, form double or triple bonds as necessary E. Exception to the octet rule 1. Duet rule for hydrogen 2. Boron usually has only six electrons about it REMEMBER BONDING PATTERNS On Some common BONDING PATTERNS C -4 bonds, N -3bonds & 1 LP, O-2 bonds & 2 LP F 1 bond & 3LP) etc

12 THE SHAPES OF MOLECULES 12.9 Predicting the Shapes of Molecules (around central atom) A. VSEPR Theory 1. Electron groups repel each other 2. Electron geometry a. Linear b. Trigonal planar c. Tetrahedral B. Predicting structures 1. Draw the Lewis structure 2. Count the number of electrons groups (Steric number) 3. Determine the number of bonding groups and lone pairs 4. Determine molecular geometry Polarity of molecules -Add up dipoles, if everything cancels, NON polar (  - Dipoles do not cancel out – Polar molecule ((  ≠  -Lone pair on central atom usually causes molecules to be polar - Unsymmetrical molecule- POLAR -MEMORIZE PARTS OF SHAPE TABLE (Bond angle and symmetry)

13 Naming Ionic Compounds A. Type I compounds 1. Metal in compounds forms only one type of ion 2. Most main group metals form type I compounds Naming type I binary ionic compounds 1. Name of cation (metal) + (base name of anion + ide) 2. Example, NaCl is sodium chloride B. Type II compounds 1. Metal forms more than one type of ion 2. Transition metals usually, but not exclusively, form type II compounds Naming type II binary ionic compounds 1. Name of cation + (charge of cation) + (base name of anion + ide) 2. Charge of cation given in roman numerals 3. FeCl 3 is iron (III) chloride C. Naming Molecular Compounds BINARY(Greek prefixes!!) A. Formed between 2 or more nonmetals B. (Prefix)name of first element + (prefix) base name of second element + ide C. 1 = mono; 2 = di; 3 = tri; 4 = tetra… D. Example, N 2 O is dinitrogen monoxide Chapter 5

14 5.5. Naming ionic compounds containing a polyatomic ion 1. Use the same procedure as ionic compounds 2. Use name of polyatomic ion, not constituent atoms 3. Example, NaNO 3 is sodium nitrate 5.6 Naming Acids A. Molecular compounds that dissolve in water to form H+ ions B. Binary acids 1. Hydrogen and nonmetal 2. (Hydro + base name of nonmetal + ic) + acid 3. Example, HCl is Hydrochloric acid C. Oxyacids 1. Hydrogen and polyatomic oxyanion 2. Oxyanions ending with –ate a. (Base name of oxyanion + ic) + acid b. Example, HNO 3 is nitric acid 3. Oxyanions ending with –ite a. (Base name of oxyanion + ous) + acid b. Example, HNO 2 is nitrous acid 5.7 Formulas and Nomenclature Summary A. Ionic compounds B. Molecular compounds C. Acids

15 EXAM POINTS 10 points Bonus question!! Part 1 Multiple Choice –Show work for partial/full credit Part 2,3 chemical names, formulas in Multiple choice Part 4 Electronic configuration, shapes of molecule, Lewis structure, Polarity Bonus question:

16 Part 1 Which of these elements have the highest electronegativity? (a) Al (b) S(c) O(d) B Part 2 Chemical formulas and write chemicals names in multiple choice Polyatomics –Match the following Part 3 Draw the Lewis structure, give the geometry of the molecule, and state whether the molecule is polar (has a dipole) or nonpolar for the following: Lewis Structure Geometry Polarity 1. CO 2 Sample Exam 2


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