Atomic Structure and the Periodic Table Chemistry – Unit 3

Early Theories of Matter Philosophers ◦ Democritus was first to propose Atomic Theory:  Matter composed of empty space through which atoms move  Atoms are indivisible ◦ Aristotle rejected Atomic Theory  Respected for ideas on nature, physics, astronomy, etc., so most ignored Democritus’ ideas

John Dalton ◦ All matter composed of atoms ◦ Atoms can not be divided ◦ Different atoms combine to form compounds ◦ Atoms are separated, combined, or rearranged in chemical reactions ◦ Conducted convincing experiments Atom – smallest particle of an element that still retains the properties of the element ◦ Can move individual atoms around to form shapes, patterns, and simple machines  Nanotechnology

Subatomic Particles and the Nuclear Atom Electron discovery ◦ Sir William Crooks discovered the cathode ray  Led to invention of TV  Cathode ray particles carry negative charge ◦ J.J. Thomson found that mass of charged particle was much less than that of a H atom  This meant Dalton was wrong and atoms are divisible  Identified the electron  Plum pudding model (p. 94 fig 4-9)

◦ Robert Millikan determined charge of electron  Single electron carries charge of -1 Nuclear atom ◦ Ernest Rutherford concluded plum pudding model was incorrect  Calculated atom consists of mostly empty space through which electrons move  Concluded there is a small, dense region in the center that contains all positive charge and virtually all mass (nucleus)  Nuclear model (p. 95 fig 4-12)

Discovery of protons and neutrons ◦ Rutherford concluded the nucleus contains positively charged particles (protons)  Protons carry a charge of +1 ◦ James Chadwick showed nucleus also contains a neutral particle in the nucleus (neutron)  Mass nearly equal to proton  Neutral charge ◦ Electrons are held within atom by attraction to positively charged nucleus ◦ Number of protons equals number of electrons

How Atoms Differ Atomic number ◦ Defined as the number of protons in an atom ◦ Determines element’s position on periodic table ◦ Atomic number = proton # = electron # Isotopes - all atoms of an element have same number of protons and electrons, but number of neutrons differ (isotopes) Mass number – sum of proton # and neutron # ◦ Number of neutrons = mass # - atomic #

Mass of individual atoms ◦ Protons and neutrons have approx. same mass ◦ Electrons are MUCH smaller ◦ B/c the masses are so small (must use scientific notation, which is cumbersome), chemists developed a standard for measurement  Carbon-12 atom  Exactly 12 atomic mass units (amu)  1 amu is 1/12 the mass of carbon-12 atom ◦ Atomic mass of an element is weighted average mass of the isotopes of that element

Unstable Nuclei and Radioactive Decay Radioactivity ◦ Chemical reactions involve only an atom’s electrons  Nucleus remains unchanged  Atom identity does not change ◦ Nuclear reactions involve change in atom’s nucleus  Atom of one element changes into atom of another element ◦ Radioactivity – some substances spontaneously emit radiation ◦ Radiation – rays and particles emitted by the radioactive material

◦ Radioactive atoms emit radiation b/c their nuclei are unstable ◦ Radioactive decay - unstable nuclei lose energy by emitting radiation spontaneously Types of radiation ◦ Experiment conducted by scientists in late 1800s determined some radiation was deflected toward positively charged plate, some toward negatively charged plate, some not at all ◦ Alpha radiation – radiation deflected toward negatively charged plate  Alpha particles  2 protons, 2 neutrons  +2 charge  Equivalent to He-4 nucleus  Represented by α

◦ Beta radiation – radiation deflected toward positively charged plate  Fast-moving electrons called beta particles  Beta particle is 1 electron with -1 charge  Represented by β ◦ Gamma rays – high-energy radiation that possess no mass and no charge  Represented by γ  B/c massless, emission of gamma rays can not result in formation of new atom ◦ Nuclear stability  Primary factor is ratio of neutrons to protons  Atoms w/ too few/many neutrons are unstable  Few radioactive atoms in nature ◦ Nuclear equation – shows atomic #, mass #, and particles involved  Both mass # and atomic # are conserved

Development of the Modern Periodic Table Modern Periodic Table ◦ Periodic law – states that there is a periodic repetition of chemical and physical properties of the elements when they are arranged by increasing atomic number ◦ Arranged in order of increasing atomic number into a series of columns, called groups, and rows, called periods ◦ Groups are labeled 1-8, followed by A and B  “A” groups are the representative elements b/c they possess a wide range of chemical and physical properties  “B” groups are the transition elements

◦ Classifying elements – 3 main classifications for elements  Metals  Located on left side of periodic table (H is exception)  Shiny, solid at room temp., good conductors, malleable, ductile  Group 1A (except H) – alkali metals  Very reactive  Group 2A – alkaline earth metals  Reactive  Transition elements  Located in the middle section of the periodic table  Transition metals  Inner transition metals – very bottom of the table to save space  Lanthanide series  Actinide series

 Nonmetals  Located in upper right side of table  Generally gases or brittle, dull solids, poor conductors  Only bromine is liquid at room temp.  Group 7A – halogens  Highly reactive  Group 8A – noble gases  Unreactive  Metalloids  Located on border of stair-step line  Physical and chemical properties of both metals and nonmetals

Classification of the Elements Organizing the elements by electron configuration ◦ Valence electrons – one of the most important relationships in chemistry  Atoms in the same group have similar chemical properties b/c they have the same number of valence electrons ◦ Valence electrons and period – the energy level of an element’s valence electrons indicates its period  Example: Li’s valence electron is in 2 nd energy level and Li is in period 2; Ga’s electron configuration is [Ar]4s 2 3d 10 4p 1, its valence electrons are in the fourth energy level, and it’s found in the 4 th period ◦ Valence electrons and group number – representative element’s group number and valence electrons are related  Noble gases have 8 valence electrons (except He)

The s-, p-, d-, and f-block elements – b/c there are 4 different energy sublevels (s, p, d, f), the periodic table is divided into 4 distinct blocks (Fig 6-10) ◦ S-block – groups 1A, 2A, and H and He ◦ P-block – groups 3A – 8A  Group 8A (noble gases) are unique b/c of their stability; they undergo virtually no chemical reactions ◦ D-block – transition metals ◦ F-block – inner transition metals ◦ Period patterns  Period 1 contains only s-block  Periods 2 and 3 contain s- and p-block  Periods 4 and 5 contain s-, p-, and d-block  Periods 6 and 7 contain s-, p-, d-, and f-block

Periodic Trends Atomic radius ◦ Trends within periods – general decrease in atomic radii left-to-right  No additional electrons come between the valence electrons and nucleus ◦ Trends within groups – general increase in atomic radii moving down  Outermost orbital increases in size along w/ increasing principal energy level, making atom larger ◦ Fig 6-12

Ionic radius ◦ Atoms can gain or lose electrons to form ions ◦ Ion – atom or bonded group of atoms that has a positive or negative charge ◦ When atoms gain electrons and form negatively charged ions, they always become larger ◦ When atoms lose electrons and form positively charged ions, they always become smaller Ionization energy – energy required to remove an electron from a gaseous atom ◦ First ionization energy - energy required to remove the first electron from an atom ◦ Indication of how strongly an atom’s nucleus holds onto its valence electrons  High ionization energy – atom has strong hold on electrons and are less likely to form positive ions

◦ Trends within periods and groups  Fig 6-17 ◦ Octet rule – atoms tend to gain, lose, or share electrons in order to acquire a full set of 8 valence electrons  Elements on right side of periodic table tend to gain electrons  Form negative ions  Elements on left side of table tend to lose electrons  Form positive ions Electronegativity – indicates the relative ability of its atoms to attract electrons in a chemical bond ◦ Fig 6-18

Properties of s-Block Elements Representative Elements ◦ The lower the ionization energy, the more reactive the metal  Metal groups – reactivity increases as the atomic number increases ◦ The higher the ionization energy, the more reactive the nonmetal  Nonmetal groups – reactivity decreases as the atomic number increases

Hydrogen ◦ Placed in group 1A only b/c it has 1 valence electron ◦ Has metallic and nonmetallic properties, so is not considered part of any group Group 1A: Alkali Metals ◦ Lose 1 valence electron and form a 1+ ion ◦ Soft ◦ Lithium  Least reactive alkali metal  Long-lasting batteries  Drug to treat bipolar disorders ◦ Sodium and potassium  Fireworks  Fertilizers

Group 2A: Alkaline Earth Metals ◦ Form compounds with oxygen (oxides) ◦ Shiny solids that are harder than alkali metals ◦ Lose 2 valence electrons to form 2+ ions ◦ Calcium  Healthy bones and teeth  Calcium carbonate  Main ingredient in limestone, chalk, and marble  Antacid tablets  Abrasives, such as toothpaste ◦ Magnesium – alloys of Mg w/ Al and Zn are strong as steel but lighter ◦ Barium – used in paints, glass; used as diagnostic tool for internal medicine

Properties of p-Block Elements Group 3A: The Boron Group ◦ Boron  Borosilicate glass for cookware  Borax - cleanser  Boric acid – disinfectant ◦ Aluminum – most abundant metal  Aluminum sulfate in anti-perspirants ◦ Gallium  Thermometers  Blue lasers

Group 4A: The Carbon Group ◦ Carbon  Organic chemistry studies C-containing compounds  Inorganic chemistry studies all others  Mineral – inorganic element found in crystals  Ore – material from which minerals can be removed  Diamond and graphite are allotropes of C  Allotropes – forms of an element in the same physical state that have different structures and properties ◦ Silicon  Semi-conductors  Sand and glass

Group 5A: The Nitrogen Group ◦ Nitrogen – 78% of earth’s atmosphere  Ammonia  TNT, nitroglycerine ◦ Phosphorus  Matchbox striking surface  Fertilizers  Fertilizers containing phosphates harm environment ◦ Arsenic – toxin used in poisons ◦ Bismuth – main ingredient in Pepto Bismol

Group 6A: The Oxygen Group ◦ Oxygen – most abundant element in earth’s crust  Bonds with most elements ◦ Sulfur  SO 2 – reacts w/water vapor to form acid rain ◦ Selenium  Vitamins  Solar panels  Photocopiers

Group 7A: The Halogens ◦ Form compounds w/ almost all metals (salts) ◦ Fluorine – most electronegative element, so greatest tendency to attract electrons  Toothpaste  Drinking water ◦ Chlorine  Disinfectant  Bleach  HCl in stomach used to digest food  PVC ◦ Iodine  Maintains healthy thyroid gland  Kills bacteria

Properties of d-Block and f-Block Elements Transition Metals ◦ Silver is best conductor ◦ Iron and titanium are used as structural materials b/c of their strength ◦ Chromium is hardest  6 unpaired electrons ◦ Magnetism – ability of a substance to be affected by a magnetic field  Moving electron creates magnetic field; b/c paired electrons move in opposite directions, their magnetic fields tend to cancel ◦ Sources of transition metals  U.S. imports more than 60 materials that are classified as “strategic and critical”

Inner Transition Metals ◦ Lanthanide series – silvery metals w/ relatively high melting points ◦ Actinide series  Radioactive  Transuranium element – atomic number >92  Plutonium is used as fuel in nuclear power plants  Americium – used in smoke detectors