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Atoms

Atoms. Simplest form of an element Composed of electrons, protons, neutrons Protons and neutrons in nucleus Electrons located outside nucleus. Protons. Positively charged Weigh approximately 1 atomic mass unit(amu). Neutrons. No charge (neutral)

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Atoms

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  1. Atoms • Simplest form of an element • Composed of electrons, protons, neutrons • Protons and neutrons in nucleus • Electrons located outside nucleus

  2. Protons • Positively charged • Weigh approximately 1 atomic mass unit(amu)

  3. Neutrons • No charge (neutral) • Weighs about the same as a proton (1 amu)

  4. Electrons • Negatively charged • About 1/2000 the size of a proton (1/2000 amu)

  5. Isotopes      Isotope- atom with same # of protons but different of neutrons.       Isotopes with more neutrons contain more mass.       Mass number – represents the sum of the protons and the neutrons in an element(always a whole number)

  6. Isotopes • Chemists often write isotopes with a specific type of notation • The element symbol(X) is written with a subscript (z) and a superscript(A) preceding it. • AZX A = the mass number, Z= the atomic number, X= the element symbol

  7. Example of Nuclear Symbol 238 = the mass number 92 = the atomic number This would be written as Uranium-238

  8. More Examples Helium-4 Lithium-7 Titanium-48

  9. Calculating Atomic Mass • Scientist developed an easier way of measuring the mass of individual atoms • Carbon-12 was assigned a mass of 12 amu’s • One atomic mass unit is the 1/12 the mass of a carbon-12 atom • Atomic mass- weighed average mass of the isotopes of that element

  10. Isotope Mass(amu) Percent abundance 6X 6.015 7.5 % 7X 7.016 92.5 % Example ISOTOPE ABUNDANCE FOR ELEMENT X Please calculate the atomic mass of element X

  11. Example Continued • Change percents to decimal number 7.5 .075 92.5 .925 • Multiply decimal number by the mass .075 x 6.015 = .45 .925 x 7.016 = 6.49 • Add the products .45 + 6.49 = 6.94 amu

  12. The Electron • Determines the behavior of most elements • 1/2000 the size of a proton • Negatively charged

  13. Basics of Quantum Mechanical Model • Electrons don’t orbit nucleus in regular patterns • Elements give off specific wavelengths of light that help identify them • Electron gaining energy= excited state • Electron losing energy=ground state

  14. Locating an Electron • Quantum numbers represent an electron’s “address” • Each number is more specific than the previous • We will concentrate on the first 3 quantum numbers

  15. 1st Quantum Number(n) • Represents the principle energy level where electron is located • Can equal a value of 1,2,3,4,5,6,or 7 • The larger the value, the farther the electron is from the nucleus • Using an analogy, this would be the “state” in which you live

  16. 2nd Quantum Number(l) • Represents the sublevel in which electron is located • Can be represented by values s,p,d,f • Each value has a different shape • Sublevel s is spherical, p is dumbbell shape; d and f too complicated to consider • For the address analogy, this would be the city in which you live

  17. 3rd Quantum Number(ml) • Represents the orientation of the electron in the sublevel • For the address analogy, this would be the street you on which you live

  18. sublevel s (sharp)

  19. Sublevel p (principle) p orbital picture

  20. Electron Configurations • Arrangement of the electrons in an atom • Aufbau principle- each electron occupies lowest energy orbital available • You must memorize this sequence

  21. Electron Configurations Continued • Sublevel s holds a max of 2 electrons • Sublevel p holds a max of 6 • Sublevel d holds a max of 10 • Sublevel f holds a max of 14 • Pauli exclusion principle – max of 2 electrons can occupy a single orientation • Hund’s Rule- electrons will fill open orbitals in the same energy level before pairing up/ Once each orbital is filled, electrons will pair but have opposite spins • Valence electrons- electrons in outermost energy level

  22. Memorizing the Pattern • Here’s a diagram to help with writing electron configurations

  23. Electrons must always enter the first available orbital of lowest energy. The first element, hydrogen, only has one electron, and so this electron must enter the 1s orbital. The electron configuration of hydrogen in the ground state must therefore be: H 1s1

  24. Pauli's exclusion principle must now be applied - the next electron to enter the 1s orbital must have a spin opposite to the spin of the electron which is already there. This completes the occupancy of the 1s orbital. The electron configuration of helium is: He 1s2.

  25. The next two electrons enter the 2s orbitals in the same way. This leads us first to the element lithium, with 3 electrons (1s2 2s1) and then, beryllium, with 4 electrons, and an electron configuration 1s2 2s2.

  26. The next element, boron, has 5 electrons. The first four of these fill the 1s and 2s levels, while the fifth enters the next level of lowest energy, which is a 2p orbital. The electron configuration of boron is: 1s2 2s2 2p1.

  27. We now apply Hund's rule: fill a set of orbitals of equivalent energy (the 2p orbitals in this case) in such a way that as many electrons as possible remain unpaired. The 6th electron enters a vacant 2p orbital rather than pairing with an unpaired electron. This element is carbon: 1s2 2s2 2p2.

  28. The next electron enters the vacant 2p orbital, giving nitrogen (1s2 2s2 2p3), oxygen (1s2 2s2 2p4) fluorine (1s2 2s2 2p5) neon (1s2 2s2 2p6). song

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