Elements may combine in more than one proportion to form more than one compound. Examples...

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1 1 UNIT 5 - ATOMIC THEORY: THE NUCLEAR MODEL OF THE ATOM 2 3 Dalton s Atomic Theory 1) Each element is made up of tiny, individual particles called atoms. 2) Atoms are indivisible; they cannot be created or destroyed. 3) All atoms of each element are identical in every respect. 4 Dalton s Atomic Theory 4) Atoms of one element are different from atoms of any other element. 5) Atoms of one element may combine with atoms of another element, usually in the ratio of small whole numbers, to form chemical compounds. Dalton s theory led to the Law of Multiple Proportions. 5 The Law of Multiple Proportions 6 Elements may combine in more than one proportion to form more than one compound. Examples... 7 Subatomic Particles Dalton s theory about the atom being indivisible was challenged with the discovery of subatomic particles within an atom. Three types we ll look at: 8 1) Electrons Electrons are negatively - charged particles that were discovered in 1820 by Michael Faraday. properties: it has a charge of -1 and a mass of x grams. 9 Electrons Thompson proposed his raisin bun model of the atom - Electrons are evenly distributed inside the spherical positive part of the atom. 10 2) Protons Protons are positively - charged particles that were discovered in 1919 by Ernest Rutherford. properties: it has a charge of +1, and has a mass that is 1836 times greater than the mass of an electron (1.67 x g). 11 3) Neutrons Neutrons are electrically neutral particles discovered in 1932 by James Chadwick. A neutron s mass is approximately equal to the mass of a proton (1.67 x grams). A summary of subatomic particles is given in Table 5.1 on page Ernest Rutherford ( ) Rutherford formed the nuclear model of the atom. Overhead: Figures 5.3 & 5.4, p. 118

2 An atom consists of a very small, positively charged nucleus, in which most of the mass of the atom is concentrated, surrounded by the electrons necessary to produce an electrically neutral atom. 13 The Nuclear Model of the Atom At the time, Rutherford only knew about electrons. Electrons circled the positively charged nucleus like the planets orbit the sun - the planetary model of the atom. 14 The Nuclear Model of the Atom Problems with Rutherford s model: A) If the electron is stationary, the atom should disappear because the electron would be electrostatically drawn in to the nucleus. B) If the electron orbits the nucleus, it will have to change direction a lot. A change in direction means acceleration, which involves the use of energy. If this is so, the electron will again spiral into the nucleus. 15 Neils Bohr ( ) Bohr was a Danish physicist ($16,000 Millionaire question!!) Bohr model All of the atom s positive charge and nearly all of its mass are contained in an extremely dense nucleus Negatively charged electrons of very small mass travel in orbits around the nucleus. 16 The Bohr model Bohr proposed that the energy of an electron in an atom cannot vary continuously, but is instead quantized - at any instant, the electron may have one of several possible energies, but at no time may it have an energy between them. 17 Where can the electrons be? The electron can orbit the nucleus at certain specified distances, but it is never found between them. example: a ladder, or steps Normally, electrons are in the Ground State - the condition when all electrons in an atom occupy the lowest possible energy levels. 18 Where can the electrons be? An electron can move between orbits by quantum jumping or quantum leaping: the electron disappears from one orbit and reappears in another. Excited State - the condition at which at least one electron in an atom is at an energy level above ground state. 19 Where can the electrons be? An electron in an excited state is unstable - will fall back to the ground state and release energy. This energy release shows up as a line in the spectrum of the element. Overhead - figure 10.3

3 20 Why was Bohr s model so good? 1) He suggested a reasonable explanation for the atomic line spectra in terms of electron energies. 2) He introduced the idea of quantized electron energy levels in the atom 21 Problems with the Bohr model 1) It only fit the hydrogen atom. 2) According to the Law of Conservation of Energy, an electron circling a nucleus should lose energy and crash into the nucleus. For the next 13 years, people worked to improve these faulty parts, leading to the Quantum Mechanical Model of the Atom. 22 Isotopes It was discovered that all atoms of an element are not identical - some atoms have more mass than other atoms of the same element. Atomic number (Z) - every atom of an element has the same number of protons. 23 Isotopes Atoms are electrically neutral, therefore the same number of electrons must be present. What s left to cause the difference in masses?? Answer: Neutrons 24 Isotopes Isotopes - atoms of the same element that have different masses (different numbers of neutrons). Mass number (A) = # protons (Z) + # neutrons The name of an isotope is its elemental name followed by its mass number. Examples: carbon, oxygen, hydrogen 25 Atomic Mass We recognize that a sample of a pure element contains atoms that have different masses The masses of atoms are expressed in atomic mass units (amu) Atomic Mass Scale handout The Periodic Table Dmitri Mendeleev and Lothar Meyer were cofounders of the periodic table in They found that certain properties repeat at regular intervals when the elements are arranged according to their atomic masses. 27 The Periodic Table However, in order for all elements to fall into the proper groups, it was necessary to switch a few of them, which interrupted the orderly increase in atomic masses. Reasons: there were errors in atomic weights as they were known in 1869 and, more importantly, it was later found that the correct ordering

4 property is the atomic number (Z), not the atomic mass. 28 The Periodic Table Horizontal rows on the table are periods. Vertical columns on the table are groups, or chemical families. Figure Period and group numbering, etc. 29 The Periodic Table Main group elements - A groups (1,2 and 13-18) Transition elements (transition metals) - B groups (3-12) Metals and nonmetals Elemental Symbols and the Periodic Table Pg. 128: Figure 5.8a - gives the name, symbol, and atomic number of the elements whose names and symbols are to be learned. Figure 5.8b - shows where they are on the periodic table. 31 Valence Electrons Are the outermost and most reactive electrons in an element. Elements in the same column, or group, of the periodic table have the same number of valence electrons, and similar chemical properties. 32 Valence Electrons Examples: Lithium, sodium, potassium Their valence shell contains only 1 electron (Group 1) All react similarly with water 33 Valence Electrons Atoms that have a full set of eight valence electrons are said to have a full octet of electrons. These atoms tend to be mostly unreactive. Example: the noble gases [Group 18 (8A)] Elements strive to attain a full octet of electrons and the same electron configuration as (become isoelectronic with) the nearest noble gas. 34 Lewis Symbols (Electron dot diagrams) The symbol of the element is surrounded by that number of dots that matches the number of valence electrons. Examples Na, C, N, Cl Note: Don t pair up electrons until there is one electron on each side of the symbol. 35 Trends in the Periodic Table ION - an atom or group of atoms that has an electrical charge because of a difference in the number of protons and electrons. eg. Na +, Cl - IONIZATION ENERGY - the energy required to remove one electron from a neutral gaseous atom. 36 Ionization Energy

5 Figure Ionization energies increase as atomic number increases within a period. Why? Larger +ve nucleus, electrons same distance away. Ionization energies decrease as atomic number increases within the group. Why? Electrons are further away. 37 Ionization Energy The amount of energy required to remove a third electron is higher than that required to remove a second electron, which is higher than the amount of energy required to remove the first electron 38 Chemical Families Elements in the same group on the periodic table have similar chemical properties. Four chemical families that we will look at: 1) the alkali metals - Group 1A (1) 2) the alkaline earths - Group 2A (2) 3) the halogens - Group 7A (17) 4) the noble gases - Group 8A (18) Figure ) The alkali metals Group 1, therefore number of valence electrons is 1. Charge on the ions? 1 +. The reactivity of the element increases as you go down the group. Why? Trends in properties of alkali metals: A) Density increases as atomic number increases. B) Boiling and melting points generally decrease as you go down the periodic table. 40 2) The alkaline earths Group 2, therefore number of valence electrons is 2. Charge on the ions? 2 +. Reactivity increases as you go down the group. 41 3) The halogens Group 17, therefore number of valence electrons is 7 (17-10). Charge on the ions? 1 -. Reactivity decreases as you go down a group. Why? The larger the atom, the less easy it is to hold onto the new electron. Density, melting point, and boiling point all increase steadily with increasing atomic number. 42 4) The noble gases

6 Group 18, therefore number of valence electrons is 8 (18-10). (Except for He - only has 2 valence electrons) Ion charges? none Are reluctant to react, because they already have a full octet of electrons, and don t want any more. Helium is fine with 2 valence electrons. Density, melting point, and boiling point all increase as you go down the group. 43 Hydrogen Valence electrons: 1. Is neither an alkali metal nor a halogen, although it shares properties with both. Can lose an electron to form H +, or gain an electron to form H Atomic Size Size of atoms increases from right to left, and from top to bottom. Figure Why? (2 reasons) 1) Distance of valence electrons from the nucleus. 2) Nuclear charge. 45 Metals and Nonmetals Figure staircase boundary Generally, an element is a metal if it can lose one or more electrons to attain a noble gas - like electron configuration Ions of metals have positive charges. Metallic character increases as you go down a group, but decreases as you go across a period from left to right. Why? Atom gets bigger, electron is further away. 46 Metals and Nonmetals As you go down a group the atom gets bigger - it s easier to lose one electron as the atom gets bigger and the electron is further away. As the period increases so does the amount of +ve charge in the nucleus - tighter hold on the valence electrons. 47 Questions: Units 5 and 10 Unit 5 - You should be able to do the target check questions and examples throughout the chapter, as well as the end - of - chapter questions that begin on pg The blue - numbered questions have answers provided at the end of the chapter. Unit 10 - questions begin on page 288 Specific questions: #3, 4, 5, 6, 8, 28, 31, 32, 33, 35, 36, 38, 39, 40, 41, 43, 44, 45, 47 (a-e, k-p), 50