Regents Chemistry Unit 1 Atomic Concepts Textbook Chapters 3 & 4
Atomic Theory- Atoms are the building blocks of matter
Atomic Models Democritus ~440 BC a Greek Philosopher suggested that matter is made up of particles so small and indestructible that they can not be divided into anything smaller John Dalton ~1797 came up with 5 principle model represented as an empty sphere 1. All matter is composed of atoms, which are invisible and indestructible. 2. All atoms of the same element are exactly alike in their chemical and physical properties. 3. All atoms of different elements are different in their chemical and physical properties. 4. Atoms of different elements combine in simple whole number ratios to form compounds 5. In chemical reactions, atoms are combined, separated or rearranged, but never created, destroyed or changed.
In 1897 an English Physicist named J.J. Thompson Came up with the Plum Pudding Model an atom. The pudding was the positive sphere and the plums were the with electrons embedded into it. The pudding is + and the plums
In 1911 Ernest Rutherford did his famous Gold Foil Experiment. Where he bombarded gold foil with alpha particles and found most of the particles passed through the gold foil and a few were deflected back.
Ernest Rutherford con t He made 2 very important conclusions from his experiment: 1) An atom has a small dense positive nucleus 2)An atom is surrounded by mostly empty space *MUST Know this!
Atomic Models (cont d) In 1913 Neils Bohr proposed a planetary model of the atom, where electrons could be found in definite orbits traveling around a positive nucleus. Solar System model of an atom. Bohr is credited for linking spectral lines to energy levels.
Atomic Models (cont d) Wave-Mechanical Model (Charge-Cloud Orbital model or Quantum Model) is presently used because electrons act like both particles and waves. It shows the most probable location of an electron in an orbit. It does this by representing the electrons as a cloud of negative charge with varying densities surrounding a small dense positively charged nucleus and is used to explain spectra of elements with multi-electron atoms.
Structure of the Atom Three subatomic particles make up the atom 1. Proton- found in the nucleus, mass =1 u, charge=+1 2. Neutron- found in the nucleus, mass =1 u, charge=0 3. Electron found outside the nucleus, mass =1/1836 the mass of a proton, charge = -1 *Nucleons- the particles found in the nucleus (protons +neutrons)
Atomic Number Atomic Number is the Number of Protons in the Nucleus this is a unique identifier for elements. No two elements have the same atomic number Is also equal to the number of electrons in an atom
Mass Number Mass Number is the number of protons and neutrons in the nucleus Mass Number is the number of nucleons This is ALWAYS a whole number Mass# - atomic# = # of neutrons
Atomic Mass Number and Atomic Mass Unit Atomic Mass (u) is defined as exactly 1/12 the mass of a carbon-12 atom; therefore 1 u is equal to the mass of carbon-12 multiplied by 1/12. Atomic Mass Number- is equal the average of the naturally occurring isotopes. NEVER expressed as a whole number
Counting Subatomic Particles Find the number of protons, neutrons and electrons in the following atoms: Na p= 11; n=12; e =11 Al p= 13; n=14; e =13 Ar p= 18; n=22; e =18
Nuclear Notation & Hyphen Notation Nuclear Notation Hyphen N otation Or C-12
Ions- charged particles Cations positive ions formed by losing electrons; ALL metals form cations Anions negative ions formed by gaining electrons; ALL non-metals form anions Find the number of protons, neutrons and electrons in the following ions: Na +1 p = 11; e = 10 ; n = 12 Cl -1 p = 17; e = 18; n= 18
Stability of Atoms Protons and Neutrons Can Form a Stable Nucleus when they have even numbers of protons and neutrons (approximately 60%) Nuclei that have mass numbers greater than 209 and atomic numbers greater than 83 are never stable
Isotopes Isotopes are the same element, same atomic number different mass number OR same element same number of protons different number of neutrons Nuclear Notation Hyphen N otation
(133 X 75/100) + (132 X 20/100) + (134 X 5/100) (133 x.75) + ( 132 x.20 ) + (134 X.05) 99.75 + 26.4 + 6.7 = 132.85 amu Isotope Calculations Ex. A sample of cesium is 75% Cs-133, 20% Cs-132 and 5% Cs-134. what is the average atomic mass? Take the mass multiplied by the corresponding percent abundance and divide by 100 plus the next mass multiplied by the corresponding percent abundance and divide by 100 plus the next mass multiplied by the corresponding percent abundance and divide by 100
Types of electrons Valence electrons outer most electrons *Use the last digit of the electron configuration on the periodic table to tell you how many! **The chemical properties of an atom are related the number of valence electrons. Inner electrons everything Except valence electrons
Lewis Dot Diagrams for atoms The valence electrons are represented using dots Hydrogen 1 electron The Lewis diagram: H Use the last digit of the electron configuration on the periodic table to tell you how many dots! You can NEVER have more than 8 dots!
Examples Oxygen - O 8 electrons to distribute; 6 valence electrons 1s 2 2s 2 2p 4 Orbital occupancy Lewis dot diagram:
Examples Oxygen - O 8 electrons to distribute; 6 valence electrons Lewis dot diagram:
Examples (cont d) Nitrogen - N 7 electrons to distribute; 5 valence electrons 1s 2 2s 2 2p 3 Orbital occupancy Lewis dot diagram:
Examples (cont d) Nitrogen - N 7 electrons to distribute; 5 valence electrons Lewis dot diagram:
Lewis Dot Diagrams for ions Negative ions (anions) are formed when an atom gains electrons. Add the necessary dot(s) with a negative charge Positive ions (cations) are formed when an atom loses electrons. No dots as the valence electrons are lost!
Electron Configuration the distribution of electrons among the various energy levels, sublevels and orbitals of an atom. They are written using principle energy level (PEL) number, the letter of the sublevel and the number of electrons in each sublevel Ex: 3p 4 An Electron Occupies the Lowest Energy Level Available (ground state) An Electron Configuration Is a Shorthand Notation
Principle Energy Levels (PELs) are used to define the energy level where electrons can be found. The principal quantum number or principle energy level (PEL) is symbolized by n principle energy levels correspond to the 7 period numbers on the periodic table
Sublevels Each PEL is divided into sublevels s sublevel has one orbital and can hold a max. of 2 electrons p sublevel has three orbitals and can hold a max. of 6 electrons d sublevel has five orbitals and can hold a max. of 10 electrons f sublevel has seven orbitals and can hold a max. of 14 electrons
Orbitals defines the region where electrons can be found No more than 2 electrons can be placed in an orbital and spin in opposite directions- Pauli Exclusion Principle Can be represented as a box with arrows indicating the electrons
Using the Periodic Table to Determine Electron Configuration 18 1 2 13 14 15 16 17 3 4 5 6 7 8 9 10 11 12
Electron Configuration Aufbau Principle shows order of filling energy levels and sublevels 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 7p 3d 4d 5d 6d 7d 4f 5f 6f 7f *Electrons are distributed from low energy levels to high energy levels Hund s Rule states that one electron must occupy an orbital before they can be paired up
Managing the Electron Hotel Floor (Energy Level) 4 1 3 5 7 Rooms 16 3 2 9 4 1 (Orbital) 1 s p d f Suites 1 electron must locate as close to the ground as possible. (1 st floor must fill before the second floor) In suites having more than 1 room (p, d & f), no room may have 2 electrons until each room has 1 electron (Hund s Rule). When rooms have 2 electrons, they must have opposite spins. (Pauli Exclusion Principle)
Ground State & Excited State Ground state electron configurations follow Auf Bau and are filled in order. Excited state electron configurations do not follow Auf Bau and have a gap, but still have the same number of electrons as a ground state configuration Ex: Carbon in the ground state: 1s 2 2s 2 2p 2 Carbon in the excited state: 1s 2 2s 1 2p 3
Ionization Energy the energy required to remove the most loosely bound electron. See Table S
An electron ion Ionization energy
Electrons and Light the electromagnetic spectrum includes all of the frequencies or wavelengths of electromagnetic radiation. It is composed of light that has a broad range of wavelengths. Our eyes can only see the visible spectrum ROYGBIV Light is an Electromagnetic Wave the electromagnetic spectrum are referred to as light even though we can not see them Bright line spectrum is the spectrum of a few colors Spectral lines - a distinct pattern of light wavelengths Light Provides Information About Electrons