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Chapter 4: Atomic Structure PowerPoint Presentation
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Chapter 4: Atomic Structure

Chapter 4: Atomic Structure

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Chapter 4: Atomic Structure

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  1. Chapter 4: Atomic Structure Atom- The basic building block of matter Learning Targets- I can describe the structure of atoms. I can describe how structure of an atom affects it’s properties. I can create a timeline that shows the developments that lead to the current model of the atom.

  2. Ancient Greek Models of atoms • 460-370 BC • Democratis- all matter consisted of extremely small particles that could not be divided he called them atomos from the Greek word Atomos which means “uncut” or “indivisiable” • 384-322 BC • Aristotle- no limit to the number of times that matter could be divided • Accepted until the 1800s

  3. The existence of atoms wasn’t scientifically proven until the early 1800s. • John Dalton, 1766-1844, English chemist and teacher • Studied the behavior of gases in air • Concluded that gas contains individual particles • No matter how large or small the sample the ratio of the elements in compounds is always the same

  4. Dalton’s atomic theory • Compounds contain atoms of more than one element • In a particular compound, atoms of different elements always combine in the same way • All elements are composed of atoms • All atoms of the same element have the same mass and atoms of different elements have different masses

  5. Daltons atomic model Tiny solid spheres with different masses Dalton’s theory explained data from many experiments and thus became widely accepted

  6. Although atoms are incredibly small they can now be observed with a scanning tunneling microscope

  7. Charged materials • Some materials when rubbed gain the ability to attract or repel other materials • Gain either a positive or negative charge • Object with like charges repel or push apart, objects with opposite charges attract or pull together • Charged particles can flow creating an electric current

  8. Thomson’s model of an atom • J.J. Thomson, an English physicist 1856-1940 • Wires connect the metal disks at opposite ends of a empty glass tube, one disk becomes negatively charged and one becomes positively charged • A glowing beam appears in the space between the plates • The beam is repelled by a negatively charged metal disk or attracted by a positively charged plates brought near it

  9. Thomson’s model • Thomson hypothesized the ray was a stream of negatively charged particles contained inside atoms, now called electrons, which are part of all atoms and carry a charge of -1. • No matter what metal he used he got the same particles • The mass was always 2000 times smaller than the mass of hydrogen atoms (a proton)

  10. Thomson’s model • First evidence that atoms are made of even smaller particles Atom is neutral- negative charges scattered throughout an atom filled with a positively charged mass of matter

  11. Ernest Rutherford’s Gold Foil Experiment • Ernest Rutherford- 1871-1937 • Atom was believed to have its positive charge spread throughout. • Rutherford shot alpha particles (large 2 + atoms) at a very thin sheet of gold foil. • If the current model of the atom was correct the alpha particles should pass though gold the mass and charge being too small to deflect the alpha particles.

  12. Rutherford’s atomic theory • Most alpha particles actually passed straight through a small fraction bounced off the gold foil at large angles • Atoms are mostly empty space • Positive charge is concentrated in the nucleus, not evenly distributed, which contains the protons and neutrons and has a positive charge • Positive charge varies among elements • Each nucleus must contain at least one proton, each proton is assigned a charge of +1

  13. Rutherford’s atomic model • All of an atoms positive charge is concentrated in the dense nucleus • Electrons are outside the nucleus

  14. James chadwick • 1932 • James Chadwick, English physicist did experiments that proved the neutrons existed • Concluded that they were neutral because they were not effected by a charged particle • Neutrons are contained in the nucleus and have a mass equal to that of a proton Compare the mass, location and charge of protons, neutrons, and electrons.

  15. Rutherford's atomic model couldn’t explain chemical properties of elements • required knowledge of electron behavior Niels Bohr (1885-1962) Focused on Electrons Agreed with Rutherford that the nucleus of and atom was surrounded by a large volume of space

  16. The Bohr Model • Electrons are only found in specific circular paths- orbits around the nucleus • Each electron orbit has a fixed energy or energy level • An electron cannot exist between energy levels • The energy level closest to the nucleus is the lowest • An electron in an atom can move from one energy level to the next when it gains or loses energy • Energy lost can be in the form of light

  17. The Quantum Mechanical Model • Rutherford-Bohr described the path of an electron as a large object would behave which was inconsistent with theoretical calculations and experimental results • Electrons move in a much less predictable way then planets in a solar system. • Erwin Schrodinger 1887-1961 • Devised and solved a mathematical equation describing the behavior of the electron in hydrogen atom

  18. Quantum Mechanical (Electron Cloud) Model • Electron Cloud- visual model of the most likely locations for electrons in an atom • Based on the probability of finding an electron with in a certain volume of space surrounding the nucleus is described as a fuzzy cloud where the electron is 90% of the time • More dense- probability high • Less dense- probability low

  19. Atomic orbitals • The electron cloud represents all the orbitals in an atom • An orbital is a region of space around the nucleus where an electron is likely to be found • Each orbital can hold 2 electrons • The lowest energy level has one orbital, the second has four, the third 9 and the fourth 16 • Electron configuration- the arrangement of electrons in the orbitals of an atom • The most stable electron configuration is the one in which the electrons are in orbitals with the lowest possible energies this is called the ground state

  20. Distinguishing Among Atoms • How are atoms of Hydrogen different than atoms of oxygen? • Element of different atoms are contain different numbers of protons

  21. Atomic Number- the number of protons in the nucleus of an atom of a given element • Identifies an Element • Atoms are electrically neutral so the number of electrons are also equal to the atomic number

  22. Elements can be represented in the following shorthand notation: • Mass Number- the number of protons and neutrons in the nucleus • Example- carbon has 6 protons and 6 neutrons so the mass number is 12 • Most the mass of an atom is concentrated in the nucleus • The number of neutrons in an atom is the difference between the mass number and the atomic number. • # neutrons = mass # - atomic # symbol Mass # Au 197 79 Atomic # Or Gold-197

  23. Isotopes- atoms that have the same number of protons but a different number of neutrons • Different mass numbers • Chemically alike- same protons and electrons which are responsible for chemical behavior

  24. Atomic Mass • The mass of atoms are givin in comparison to carbon-12 • An atomic mass unit (amu) = 1/12 the mass of a carbon-12 atom • Carbon 12 amu • Flourine- actual mass 3.155 x 10 –23 g, atomic mass- 18.998 amu • 1 proton or neutron- 1 amu • Atomic mass- weighted average mass and relative abundance of isotopes as they occur in nature • Example- Hydrogen –1, 99 %, 1.0078 • Hydrogen- 2, 1% heaver • Hydrogen- 3 Atomic mass- 1.0079