SCI1055 Chapter 8 PDF

2.1 Elements (1) An element is a pure substance that cannot be broken down into simpler substances by a chemical reaction. Each element is identified by a one- or two-letter symbol. Elements are arranged in the periodic table. The position of an element in the periodic table tells us much about...

2.1 Elements (1) An element is a pure substance that cannot be broken down into simpler substances by a chemical reaction. Each element is identified by a one- or two-letter symbol. Elements are arranged in the periodic table. The position of an element in the periodic table tells us much about its chemical properties. 1  2.1 Elements (2) Table 2.1 Common Elements and Their Symbols Element Symbol Element Symbol Bromine Br Magnesium Mg Calcium Ca Manganese Mn Carbon C Molybdenum Mo Chlorine Cl Nitrogen N Chromium Cr Oxygen O Cobalt Co Phosphorus P Copper Cu Potassium K Fluorine F Sodium Na Hydrogen H Sulfur S Iodine I Zinc Zn Lead Pb 2 2.1 Elements (3) 3  2.1 Elements (4) A. Elements and the Periodic Table The elements in the periodic table are divided into three groups—metals, nonmetals, and metalloids. Metals: They are located on the left side of the periodic table. They are good conductors of heat and electricity. Metals are shiny solids at room temperature, except for mercury (Hg), which is a liquid. 4  2.1 Elements (5) A. Elements and the Periodic Table Nonmetals: They are located on the right side of the periodic table. Nonmetals have a dull appearance They are usually poor conductors of heat and electricity. Nonmetals can be solids, liquids, or gases at room temperature solid Examples: sulfur, carbon liquid Examples: bromine gas Examples: nitrogen, oxygen 5  2.1 Elements (6) A. Elements and the Periodic Table Metalloids: These are located on the solid line that starts at boron (B) and angles down towards astatine (At). Metalloids have properties intermediate between metals and nonmetals Only seven elements are Metalloids: boron (B) antimony (Sb) silicon (Si) tellurium (Te) germanium (Ge) astatine (At) arsenic (As) 6  2.1 Elements (10) C. Compounds Compound: a pure substance formed by chemically combining two or more elements together. A chemical formula consists of: Element symbols to show the identity of the elements forming a compound. Subscripts to show the ratio of atoms in the compound. H2O C3H8 2 H atoms 1 O atom 3 C atoms 8 H atoms 7  2.2 Structure of the Atom (1) All matter is composed of the same basic building blocks called atoms. Atoms are composed of three subatomic particles: Table 2.3 Summary: The Properties of the Three Subatomic Particles Subatomic Particle Charge Mass (g) Mass (amu) Proton +1 1.6726 10 24 1 Neutron 0 1.6749 10 24 1 Electron −1 9.1093 10 28 Negligible 8  2.2 Structure of the Atom (2) Nucleus: Electron cloud: location of protons location of electrons and neutrons comprises most of the dense core of the atom atom’s volume location of most of the mostly empty space atom’s mass 9  2.2 Structure of the Atom (4) From the periodic table: 3 Atomic number (Z) is the number of protons Li in the nucleus. Every atom of a given element has the same number of protons in the nucleus. Different elements have different atomic numbers. A neutral atom has no net overall charge, so Z = number of protons = number of electrons 10  2.3 Isotopes (1) 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) Mass number (A) = + the number of neutrons Mass number 35 Cl 37 Mass number (A) 17 17 Cl (A) Atomic number (Z) Atomic number (Z) # of protons = 17 # of protons = 17 # of electrons = 17 # of electrons = 17 # of neutrons = 35 − 17 = 18 # of neutrons = 37 − 17 = 20 11  2.3 Isotopes (2) 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 atomic number C element symbol 12.01 atomic weight (amu) 12  2.3 Isotopes (3) B. Atomic Weight HOW TO Determine the Atomic Weight of an Element Example What is the atomic weight of chlorine? List each isotope, its mass in atomic Step 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 13  2.3 Isotopes (4) B. Atomic Weight HOW TO Determine the Atomic Weight of an Element Step Multiply the isotopic abundance by the mass of each isotope, and add up the products. The sum is the atomic weight of the element. 34.97 × 0.7578 = 26.5003 amu 36.97 × 0.2422 = 8.9541 amu 4 sig. figs. 35.4544 amu = 35.45 amu 14  2.4 The Periodic Table (1) A. Basic Features of the Periodic Table A row in the periodic table is called a period, and a column in the periodic table is called a group. Main group elements: They consist of the tall columns on the right and left of the Periodic Table. The groups are numbered 1A to 8A. Transition metal elements: These are in the 10 short columns in the middle. The groups are numbered 1B to 8B. Inner transition elements: They consist of the lanthanides and actinides. 15 There are no group numbers assigned.  2.4 The Periodic Table (2) A. Basic Features of the Periodic Table 16  2.4 The Periodic Table (3) B-1 Characteristics of Groups 1A and 2A Elements that comprise a particular group have similar chemical properties. Properties of Both Group Group Groups Number Name soft and shiny metals 1A Alkali metals low melting points 2A Alkaline earth elements good conductors of heat and electricity react with water to form basic solutions 17  2.4 The Periodic Table (4) B-2 Characteristics of Groups 7A and 8A Group Group Number Name Properties exist as two atoms joined together 7A Halogens very reactive very stable 8A Noble gases rarely combine with any other elements 18  Classical “atoms” Predictions of classical theory Electrons orbit the nucleus Curved path = acceleration Accelerated charges radiate Electrons lose energy and spiral into nucleus Atoms cannot exist! Experiment - atoms do exist  New theory needed Atomic spectra Blackbody radiation Continuous radiation distribution Depends on temperature of radiating object Characteristic of solids, liquids and dense gases Line spectrum Emission at characteristic frequencies Diffuse matter: incandescent gases Illustration: Balmer series of hydrogen lines  The quantum concept Max Planck (1900) Introduced quantized energy Einstein (1905) Light made up of quantized photons Higher frequency h 6.63 10 34 Joule sec ond photons = more energetic photons  Bohr’s theory Three rules: Electrons only exist in certain allowed orbits Within an orbit, the electron does not radiate Radiation is emitted or absorbed when changing orbits (quantum leaps)  Quantum theory of the atom Lowest energy state = “ground state” Higher states = “excited states” Photon energy equals difference in state energies Hydrogen atom example Energy levels Line spectra Photon energy: hf EH  EL  Quantum mechanics Bohr theory only modeled the line spectrum of hydrogen Did not work for atoms larger than hydrogen New, better theory needed Further experiments established wave-particle duality of light and matter Light has both wave and particle properties  Wave Particle Duality Louis de Broglie (1923) Postulated that if a particle of light has a dual nature, then particles such as electrons should also Electrons confined to space near nucleus, therefore must be confined (standing) waves Confined waves Only certain fundamental frequencies and harmonics exist Pattern depends on wavelength and velocity New theory – wave (quantum) mechanics  The quantum mechanics model Highly mathematical treatment of matter waves Electron considered as a spread-out wave Three dimensional Knowledge of electron location is uncertain Heisenberg Uncertainty Principle The position and momentum of electron cannot be measured Location described in terms of probabilities Orbital - Fuzzy region of space where electron is likely to be found Characteristic 3-D shapes (Probability cloud) Identified with characteristic energy levels Quantum numbers specify electronic quantum states Electronic quantum numbers in atoms Principle quantum number, n Energy level Average distance from nucleus Angular momentum quantum number, l Spatial distribution Labeled s, p, d, f, g, h, … Magnetic quantum number Spatial orientation of orbit Spin quantum number Electron spin orientation  2.5 Electronic Structure (1) An electron is confined to a specific region around the nucleus, giving it a particular energy. The regions occupied by electrons are called principal energy levels or shells (n). The shells are numbered n = 1,2,3,4 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. 28  2.5 Electronic Structure (2) Shells with larger numbers (n) are farther from the nucleus and can hold more electrons. The maximum number of electrons in each shell is given by the formula 2n 2 , where n = shell number. The distribution of electrons in the first four shells: Number of Electrons Shell (n) in a Shell 4 32 3 18 increasing increasing 2 8 number of energy electrons 1 2 29  2.5 Electronic Structure (3) 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. Subshell Number of Orbitals s 1 increasing p 3 energy d 5 f 7 30  2.5 Electronic Structure (4) Table 2.4 Orbitals and Electrons Contained in the Principal Energy Levels (n = 1 − 4) Electrons in Each Maximum Number Shell Orbitals Subshell of Electrons 1 1s 2 2 2 2s 2 8 2p 2p 2p 3×2=6 3 3s 2 18 3p 3p 3p 3×2=6 3d 3d 3d 3d 3d 5 × 2 = 10 4 4s 2 32 4p 4p 4p 3×2=6 4d 4d 4d 4d 4d 5 × 2 = 10 4f 4f 4f 4f 4f 4f 4f 7 × 2 = 14 31  2.5 Electronic Structure (5) The s orbital has a spherical shape. The orbital gets larger in size as the shell number increases. The p orbital has a dumbbell shape. 32  2.6 Electron Configuration (1) 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 Electrons are placed in the lowest energy orbital beginning with the 1s orbital. Orbitals are then filled in order of increasing energy. 33  2.6 Electron Configuration (2) Rules to Determine the Ground State Electronic Configuration of an Atom 34 Electron Configurations and the Periodic Table 35  Electron Configuration Rules to Determine the Ground State Electronic Configuration of an Atom Rule Each orbital holds a maximum of 2 electrons. Rule 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. 36  2.6 Electron Configuration (6) 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. Electron Noble Gas Configuration Notation element: C 1s 2 2 s 2 2 p 2 [He] 2 s 2 2 p 2 nearest He 1s 2 noble gas: 37  2.7 Valence Electrons (1) 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 1s 2 2 s 2 1s 2 2 s 2 2 p 6 3 s 2 3 p 5 valence shell: n = 2 valence shell: n = 3 valence electrons = 2 valence electrons = 7 38  2.7 Valence Electrons (2) 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 to 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. 39  2.7 Valence Electrons (3) A. Relating Valence Electrons to Group Number Group number: 1A 2A 3A 4A 5A 6A 7A 8A Period 1: H He 1s1 1s 2 Period 2: Li Be B C N O F Ne 2s1 2s 2 2 s 2 2 p1 2 s 2 2 p 2 2 s 2 2 p 3 2 s 2 2 p 4 2 s 2 2 p 5 2 s 2 2 p 6 Period 3: Na Mg Al Si P S Cl Ar 2 2 2 4 2 6 3s1 3s 2 3 s 2 3 p1 3 s 3 p 3 s 2 3 p 3 3 s 3 p 3 s 2 3 p 5 3 s 3 p 40