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Section 4.3

Section 4.3. How Atoms Differ. Objectives. Explain the role of atomic number in determining the identity of an atom. Define an isotope and explain why atomic masses are not whole numbers.

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Section 4.3

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  1. Section 4.3 How Atoms Differ

  2. Objectives • Explain the role of atomic number in determining the identity of an atom. • Define an isotope and explain why atomic masses are not whole numbers. • Calculate the number of electrons, protons, and neutrons in an atom given its mass number and atomic number.

  3. Atomic Number • Moseley (1887-1915) discovered that atoms ofeach element contain a unique number of positive charges in their nuclei. • Therefore, it was concluded that the number of protons identifies an atom as an atom of a particular element. • The number of protons in an atom is referred to as an element’s atomic number.

  4. Atomic Number • Atomic number determines an element’s position in the periodic table - from left to right and top to bottom, elements are arranged by increasing atomic number. Carbon Element name ATOMIC NUMBER Chemical symbol Average atomic mass 6 C 12.011

  5. Atomic Number • Recall that atoms are neutral. • The number of protons, therefore, MUST be equal to the number of electrons. • Atomic number = number of protons = number of electrons • Carbon, for example, has an atomic number of 6 and 6 protons and 6 electrons.

  6. Complete the following table. Practice Problems

  7. Isotopes • Recall that Dalton was wrong about atoms being indivisible. He was also wrong when he stated that all atoms of an element are identical. • All atoms of an element DO have the same number of protons and electrons. The number of neutrons, however, MAY DIFFER from atom to atom. • Atoms that have the same number of protons but different numbers of neutrons are called isotopes.

  8. 3 Isotopes of Hydrogen • Each has 1 proton. • They differ in the number of neutrons they contain. • Because of this, each atom has a different mass. • Note that the isotopes of hydrogen have special names.

  9. Isotopes • Elements are found in nature as a mixture of isotopes. • The percent abundance of the isotopes (percent of each isotope in a sample) is constant no matter where the sample is found. • Isotopes have the same chemical behavior because chemical behavior is determined by the number of ELECTRONS an element has.

  10. Isotopes and Mass Number • To make it easy to identify the various isotopes of an element, scientists add a number after the element’s name. • This number is the mass number - the sum of the number of protons and neutrons in the nucleus of an atom of the isotope. • Mass # = # protons + # neutrons

  11. Example • Oxygen-17 • oxygen has an atomic number of 8, so oxygen-17 has 8 protons and electrons. • Since mass # = # protons + # neutrons, then # neutrons = mass # - #protons. So, in oxygen-17, # neutrons = 17 - 8 or 9. • Oxygen-17 has 8 protons, 8 electrons, and 9 neutrons.

  12. Mass number Isotopes Atomic number • Chemists use another shortened notation to identify isotopes. • The chemical symbol for the element is used. Written to the left of the symbol, a superscript indicates the mass number and a subscript indicates the atomic number.

  13. Practice Problems

  14. Atomic Mass Units • Recall that the masses of protons, neutrons, and electrons are VERY small numbers. These numbers are difficult to work with. • Scientists therefore developed a way to measure their mass relative to a standard atom (the carbon-12 atom). • One atomic mass unit(amu) is defined as 1/12 the mass of a carbon-12 atom.

  15. Atomic Mass Units • The masses of the particles in terms of amu are: • electron = 0.000549 amu (We round to 0 amu) • proton = 1.007276 amu (We round to 1 amu) • neutron = 1.008665 amu (We round to 1 amu) • Note that the mass of a proton & neutron is slightly different; also, correct masses of atoms will not be whole numbers but decimals.

  16. Atomic Mass • The atomic mass of an element is the weighted average mass of the isotopes of that element. • Atomic mass, therefore, takes into account the percent abundance of each isotope as well as the mass of an atom of that isotope. • The isotope in greater abundance will have a greater effect on determining the atomic mass.

  17. Analyzing Atomic Mass • Analyzing an element’s atomic mass can SOMETIMES give you a clue as to what the most abundant isotope for that element is. • The atomic mass of fluorine is 18.998 amu. From this you can conclude that most fluorine occurs in the form of fluorine-19.

  18. Calculating Atomic Mass • Multiply the % abundance of each isotope by its mass. 34.969(75.77%) = 26.496 amu 36.966(24.23%) = 8.9569 amu • Add the mass contributions of all the isotopes together. 35.453 amu is the atomic mass of chlorine. Chlorine-35 Chlorine-37 Atomic mass: 34.969 amu Atomic mass: 36.966 amu % abundance: 75.77% % abundance: 24.23%

  19. Practice Problems • Use the following information to calculate the atomic mass of magnesium: magnesium-24: atomic mass 23.985, % abundance 78.99% magnesium-25: atomic mass 24.986, % abundance 10.00% magnesium-26: atomic mass 25.982, % abundance 11.01% • Use the following information to calculate the atomic mass of X; then identify X. 6X = 6.015 amu; 7.5% abundance 7X = 7.016 amu; 92.5% abundance

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