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Chemistry Unit 3 Atomic Structure (Ch.3) 3-1 Early Models of Atoms Democritus (450 BC) Proposed that all matter was made of tiny indivisible particles. He called these particles atomos (meaning indivisible). We call them atoms. Good looking guy! Atom An atom is the smallest particle of an element that retains the identity of that element. If we repeatedly cut a piece of Al, the smallest possible piece is an atom of Al. Classic model of an atom Aristotle Didn’t agree with Democritus. Believed matter was continuous and made up of only one substance called “hyle” It wasn’t until the 1700’s when his ideas were reexamined. Newton and Boyle (1600s) Published articles stating their belief in the atomic nature of elements They had no proof Antoine Lavoisier (1770’s) French, The “Father of Modern Chemistry” Discovered the law of conservation of matter. Matter is neither created nor destroyed. Joseph Proust (1799) French Chemist Developed The Law of Definite Proportions Compounds always contain elements in the same proportion by mass. Law of Definite Proportions H20 (by mass is always) 88.9% Oxygen, 11.1% Hydrogen If we had an 80g sample of H20 how much is O? .889 x 80 = 71g How much is H? .111 x 80 = 9g John Dalton (1803) Proposed the atomic theory of matter, which is the basis for present atomic theory John Dalton, English schoolteacher Atomic Theory of Matter Each element is composed of extremely small particles called atoms. All atoms of a given element are identical, but differ from those of any other element. Which element is this? Atomic theory of matter When elements unite to form compounds, they do so in a ratio of small whole numbers. This is called the Law of Multiple Proportions. Ex: C and O can combine to form CO or CO2, but not CO1.8. Dalton’s Model of an Atom All matter is composed of tiny particles J.L. Gay-Lussac (early 1800s) Observed that working with gas reactions at constant volume, temperature and pressure are directly related. He named the discovery of this relationship Charles Law, which is represented by P1/T1=P2/T2. Amadeo Avogadro (early 1800s) – Italian Physicist Explained GayLussac’s work using Dalton’s theory: Equal volumes of gases at the same temp/pres have the same number of gas molecules. Michael Faraday (1839) Suggested that atomic structure was related to electricity. Atoms contain particles that have electrical charges. Positive (+) Negative (-) Opposite charges attract Like charges repel William Crookes (1870’s) English Physicist Developed the cathode ray tube to find evidence for the existence of particles within the atom. J.J. Thomson (1896) Used a cathode ray tube (CRT) to identify negatively charged particles, called electrons. Determined the ratio of an electron’s charge to its mass. Developed the “plum pudding” model of an atom. Cathode ray bending toward a positive charge Plum Pudding Model - - + + + + + + - + Atoms are composed of randomly arranged charged particles Robert Millikan (early 1900s) US Physicist Used the oil drop experiment to prove the electron has a negative charge Was able to determine the charge of the electron Millikan’s Oil Drop Experiment Bothe/Chadwick (early 1930s) English Found high energy particles with no charge with the same mass as the proton called neutrons. Ernest Rutherford (1909) Used the gold foil experiment to prove the atom is mostly empty space. Rutherford concluded that all of an atom’s positive charge, and most of its mass is located in the center, called the nucleus. Analogy: thumb nail and the 50 yard line. 98% of the particles passes straight through 2% of the particles deflected off at varying angles 0.01% of the particles bounced straight back Rutherford’s Planetary Model of an atom Positive charge and majority of mass located in the nucleus. Negatively charged electrons orbit the nucleus like planets. - ++++ ++ + - - - Most of an atom is empty space! Problem He thought a moving electrical charge (-) in a curved path should lose energy (give off light). If it did, it would fall into the (+) nucleus. Why don’t the (-) electrons fall into the (+) nucleus? Atom:The smallest particle of an element that has the properties of that element. Make up of nucleus consists of protons and neutrons Surrounded by an electron cloud Electron cloud Sub-Atomic Particles Protons Positively (+) charged The number of protons in an atom refers to the atomic number (Z) Composed of 3 quarks (2 up, 1 down) Mass= 1.6726 x 10-27kg Atomic mass 1 amu (µ) Sub-Atomic Particles Neutrons Found in nucleus Neutral (no) charge composed of 3 quarks (1 up, 2 down) Atomic mass 1 amu (µ) Isotopes- atoms of the same element that have a different number of neutrons. Sub-Atomic Particles Electrons Found in electron clouds surrounding the nucleus. Negative (-) charge Mass = .00091 x 10-27 kg 1800 times smaller than protons & neutrons Mass 0 amu (µ) Sub-atomic particles Electrons Orbit the nucleus at very high speed in energy levels (electron clouds). Negatively (-) charged Have no mass (when compared to protons and neutrons) Atomic Number = Protons The atomic number of an element is the number of protons an element has. Located above the symbol of the element The number of protons determines the identity of the element. Each element has a different atomic number Neils Bohr (1913) Improved Rutherford’s work by saying electrons do not lose energy in the atoms so they will stay in orbit Stated there are definite levels in which the electrons follow set paths without gaining or losing energy (Planetary Model) Each level has a certain amount of energy associated with it and the electrons can only jump levels if they gain or lose energy Could not explain why (-) electrons don’t fall into the (+) nucleus. Energy Levels In the ground state for an atom, electrons are at their lowest, most stable energy levels. In the excited state, atoms require energy and electrons move to a higher energy level. How many electrons are in an atom? For an atom to have an overall neutral charge the number of electrons must equal the number of protons. #Protons=#electrons What element is this? Mass number The Mass number of an atom is the sum of the mass of protons and neutrons Located below the symbol of the element Atomic mass is measured in amu’s, (atomic mass units) Based on Carbon having a mass of 12 Mass = Protons + Neutrons How many neutrons are in an atom? Mass=Protons+Neutrons 195= 78 + Neutrons 195-78= Neutrons Platinum has 117 Neutrons Find the number of neutrons in: Hydrogen Carbon Helium Potassium Boron Gold Mass =Protons + Neutrons Hydrogen (H) 1 =1 + Neutrons Hydrogen has 0 neutrons Helium (He) 4 = 2 + Neutrons Helium has 2 neutrons Boron (B) 11 = 5 + Neutrons Boron has 6 neutrons Carbon (C) 12 = 6 + Neutrons Carbon has 6 neutrons Potassium (K) 39 = 19 + N Potassium has 20 neutrons Gold (Au) 197 = 79 + N Gold has 118 neutrons Atomic Mass The average mass of all of the isotopes of an element. Aka: average atomic mass number, or atomic weight. Isotopes:Atoms of the same element with different masses. Average Atomic Mass Ne-20 has a mass of 19.992 amu (u), and Ne-22 has a mass of 21.991 amu (u). In any sample of 100 Ne atoms, 90 will be Ne-20. Calculate the average atomic mass of Ne. .90 x 19.992 = 17.9928 .10 x 21.991 = 2.1991 avg mass = 20.1919 amu Ions An atom that has gained or lost an electron. It acquires a net electrical charge. If an atom loses an electron (oxidation) it has more protons than electrons and has a net positive charge. (cation) Na 11 P 11 e- 11 P 10 e- Na+ Ions If an atom gains an electron (reduction) it has more electrons than protons and has a net negative charge.(anion) 7 valence e- Full octet Ionic Charges Charge of ion = # protons - # electrons What is the charge of a magnesium atom that loses 2 electrons? Number of protons 12 -Number of electrons 10 charge of ion +2 Mg2+ or Mg+2 Charge is written to the upper right of the symbol. Representations of atoms General form: (Elemental Notation) X = Element Symbol A = Atomic Mass (P + N) Z = Atomic Number (P) Ionic Charge A Z Charge X What is the atomic structure? Determine the number of: P = N = e = 23 + Na 11 What is the atomic structure? Determine the number of: P = 11 N = 12 e = 10 23 + Na 11 What is the atomic structure? Determine the number of: P = N = e = 127 I 53 - What is the atomic structure? Determine the number of: P = 53 N = 74 e = 54 127 I 53 - Put into elemental notation Atomic # = 29 Atomic Mass = 64 Ionic charge = +2 ? How many electrons? Atomic # = 29 Atomic Mass = 64 Ionic charge = +2 # of electrons = 64 2+ 29 Cu Put into elemental notation 37 Protons 48 Neutrons 36 Electrons ? Put into elemental notation 37 Protons 48 Neutrons 36 Electrons 85 37 + Rb Max Planck (early 1900s) Proposed Planck’s Theory which says that energy is given off in little packets or particles called quanta which is based on the particle nature of light Each quantum of energy corresponds to the different energy levels for electrons. Proposed the equation: E=hf, where E is energy, f is frequency, and h is Planck’s constant (6.63 x 10^-34 J/Hz) De Broglie (1923) Suggested that if waves can have a particle nature, particles can have a wave nature, known as the “waveparticle duality” principle Wondered why the positive nucleus and negative electrons do not attract. Proposed that electrons moved so fast (speed of light) that they had properties of waves instead of particles. The Study of Waves Wave: a progressive disturbance propagated from point to point in a medium or space without progress or advance by the points themselves Types of Waves Mechanical: a wave that requires an energy source and an elastic material medium to travel. Electromagnetic: a wave that does not require a material medium to travel; it propagates by electric and magnetic fields Wave Travel Transverse: displacement of the medium is perpendicular to the direction of propagation of the wave. Longitudinal: displacement of the medium is parallel to the direction of propagation of the wave Properties of Waves Wavelength ()ג: The distance between any part of the wave (peak) and the nearest part that is in phase with it (another peak). Standard unit is meters (m). Frequency (f ): The number of peaks which pass a given point each second. Standard unit is cycles per second which is a hertz (Hz). Amplitude (A): The maximum displacement of a vibrating particle from its equilibrium position. Standard unit is meters (m). Velocity (v): the distance a wave (peak) travels in a given time. Standard unit is meters per second (m/s). Energy (E): The energy of a single photon of radiation of a given frequency. Standard unit is the joule (J). Some relationships between the properties of waves are represented by the equations: V=f* גand E=h*f , where h=6.63x10^-34 J/Hz Werner Heisenberg (1927) Proposed his “Heisenberg uncertainty principle”, which says that the position and momentum of an electron cannot simultaneously be measured and known exactly. The arrangement of electrons is discussed in terms of the probability of finding an electron in a certain location. Erwin Schrodinger (1926) Studied the wave nature of the electron and developed mathematical equations to describe their wave-like behavior. The most probable location of the electrons can be found and the plot of this probability is called the charge cloud model. The four quantum numbers Principal Quantum Number (n) Refers to the energy levels in the atom which is the distance from the nucleus and designated with a positive whole number (n=1,2,3,etc) Wavelength of emitted photon is determined by the “energy jump” between energy levels Energy levels (or shells) means electrons are contained in an area where the probability of finding the electron is 90% Angular Momentum Quantum Number (l ) Refers to the sublevel (within an energy level) which is one or more “partitions” each with a slightly different energy. The number of sublevels in a particular energy level is equal to the principal quantum number (n). The types of sublevels include: s, p, d, f, etc. The four quantum numbers (continued) Magnetic Quantum Number (m) Refers to the orientation in space of the electrons in a sublevel Can have any whole number value from -1 to +1 which will tell how many orbitals are in a sublevel. A maximum of 2 electrons per orbital. Sublevel s p d f # of Orbitals Total # of electrons Four Quantum Numbers (continued) Spin Quantum Number (s) Refers to the spin of the electron. Pauli Exclusion Principle : if two electrons occupy the same orbital, they must have opposite spin. Half-filled orbital: _____ Filled orbital: _____ Permissible Values of Quantum Numbers for Atomic Orbitals n 1 l 0 m 0 Orbital 1s # of Subshells 1 #of Orbitals 1 Max # of Electrons 2 2 0 1 0 -1,0,1 2s 2p 2 1 3 2 6 3 0 1 2 0 -1,0,1 -2,-1,0,1,2 3s 3p 3d 3 1 3 5 2 6 10 4 0 0 4s 1 -1,0,1 4p 2 -2,-1,0,1,2 4d 3 -3,-2,-1,0,1,2,3 4f 4 1 3 5 7 2 6 10 14 Distribution of Electrons for Different Elements (Electron Configuration) Electrons will occupy the lowest energy levels and sublevels first. Notation: Principal Quantum Number, n (energy level) 2p Type of Orbital (sublevel) Number of electrons 2 y Orientation of Orbital Long Notation: Pyramid Filling “Rule of thumb” for filling electrons at the lowest energy level possible. Give the long notation electron configuration for: O Ca Ag Give the short notation O Ca Ag Orbital Diagrams Usually only done for the outer shell electrons, which always includes the s and p orbitals. Electron Dot Diagrams Shows the outer shell electrons for elements.