Chemistry - Chp 4 - Atomic Structure - Notes
- 1. Chapter 4 Atomic Structure
- 2. Section 4.1 Defining the Atom Objectives: • Describe Democritus’s ideas about atoms • Explain Dalton’s atomic theory • Identify what instrument is used to observe individual atoms Defining the Atom • The Greek philosopher Democritus (460 B.C. – 370 B.C.) was among the first to suggest the existence of atoms (from the Greek word “atomos”) • He believed that atoms were __________________________ and ____________________________ • His ideas did agree with later scientific theory, but did not explain chemical behavior, and was _____________________________________________________________ - but just philosophy Dalton’s Atomic Theory 1. All elements are composed of tiny indivisible particles called ____________________ 2. Atoms of the same element are _____________________________. Atoms of any one element are different from those of any other element. 3. Atoms of different elements ________________________________ in simple whole number ratios to form _____________________________________________________ 4. In chemical reactions, atoms are ___________________________________________ ___________________________________________________________ - but never changed into atoms of another element. Sizing up the Atom • Elements are able to be subdivided into smaller and smaller particles – these are the atoms, and they still have properties of that element • If you could line up 100,000,000 copper atoms in a single file, they would be approximately 1 cm long 2
- 3. • Despite their small size, individual atoms are observable with instruments such as ____________________________________________________________________ _ Section 4.2 Structure of the Nuclear Atom Objectives: • Identify three types of subatomic particles • Describe the structure of atoms, according to the Rutherford atomic model Structure of the Nuclear Atom • One change to Dalton’s atomic theory is that ________________________________ into subatomic particles: ex.) • • • Discovery of the Electron • In 1897, J.J. Thomson used a _______________________________ to deduce the presence of a negatively charged particle: _________________________________ • Cathode ray tubes pass electricity through a gas that is contained at a very low pressure. Mass of the Electron • 1916 – Robert Millikan determines the mass of the electron: _________________ the mass of a hydrogen atom; has one unit of negative charge Conclusions from the Study of the Electron: a) Cathode rays have identical properties regardless of the element used to produce them. All elements must contain identically charged electrons. b) Atoms are neutral, so there must be ___________________________________ in the atom to balance the negative charge of the electrons 3
- 4. c) ____________________________________________________ that atoms must contain other particles that account for most of the mass • Eugen Goldstein in 1886 observed what is now called the ________________- particles with a positive charge, and a relative mass of 1 (or 1840 times that of an electron) • 1932 – James Chadwick confirmed the existence of the __________________ – a particle with no charge, but a mass nearly equal to a proton Subatomic Particles Particle Charge Mass (g) Location Electron (e-) Proton (p+) Neutron (no) Thomson’s Atomic Model • Thomson believed that the electrons were like plums embedded in a positively charged “pudding,” thus it was called the “plum pudding” model. 4
- 5. Ernest Rutherford’s Gold Foil Experiment - 1911 • Alpha particles are helium nuclei - The alpha particles were fired at a thin sheet of gold foil • Particles that hit on the detecting screen (film) are recorded Rutherford’s Findings • Most of the particles passed right through • A few particles were deflected • VERY FEW were greatly deflected Conclusions: a) b) c) The Rutherford Atomic Model • Based on his experimental evidence: • The atom is mostly empty space • All the positive charge, and almost all the mass is concentrated in a small area in the center. He called this a _____________________________ 5
- 6. • The nucleus is composed of _________________________ and ______________________________ (they make the nucleus!) • The _________________________________ distributed around the nucleus, and occupy most of the volume • His model was called a ___________________________________ 6
- 7. Name _____________________________________________ Date ________________ 4.1 and 4.2 Section Review 1. In your own words, state the main ideas of Dalton’s atomic theory. 2. Democritus and Dalton both proposed that matter consists of atoms. How did their approaches to reaching that conclusion differ? 3. What are the charges and relative masses of three main subatomic particles? 4. Describe the basic structure of an atom 5. Describe Thomson’s, Millikan’s, and Rutherford’s contributions to atomic theory. Include their experiments if appropriate 7
- 8. Section 4.3 Distinguishing Among Atoms Objectives: • Explain what makes elements and isotopes different from each other • Calculate the number of neutrons in an atom • Calculate the atomic mass of an element • Explain why chemists use the periodic table Atomic Number • Atoms are composed of identical protons, neutrons, and electrons o How then are atoms of one element different from another element? • Elements are different because they contain different numbers of _______________________ • The _________________________________ of an element is the ____________________________________________ in the nucleus • # protons in an atom = __________________________________________ • Atomic number (Z) of an element is the number of protons in the nucleus of each atom of that element. Element # of protons Atomic # (Z) Carbon Phosphorus 8
- 9. Gold Mass Number • Mass number is the number of protons and neutrons in the nucleus of an isotope: Nuclide p+ n0 e- Mass # Complete Symbols • Contain the symbol of the element, the mass number and the atomic number. 9
- 10. Isotopes • Dalton was wrong about all elements of the same type being identical • Atoms of the same element can have different numbers of ______________________ • Thus, different mass numbers. • These are called _________________________________ • Frederick Soddy (1877-1956) proposed the idea of isotopes in 1912 • ____________________________ are atoms of the ______________________________________ having different masses, due to varying numbers of neutrons. • Soddy won the Nobel Prize in Chemistry in 1921 for his work with isotopes and radioactive materials. Naming Isotopes • We can also put the mass number after the name of the element: o carbon-12 o carbon-14 o uranium-235 Isotope Protons Electrons Neutrons Nucleus Hydrogen – 1 1 1 0 (protium) Hydrogen – 2 1 1 1 (deuterium) 10
- 11. Hydrogen – 3 1 1 2 (tritium) Isotopes • Elements occur in nature as _________________________ of isotopes. • Isotopes are atoms of the same element that differ in the number of neutrons. Atomic Mass • How heavy is an atom of oxygen? o It depends, because there are different kinds of oxygen atoms. • We are more concerned with the __________________________________________ • This is based on the abundance (percentage) of each variety of that element in nature. o We don’t use grams for this mass because the numbers would be too small. Measuring Atomic Mass • Instead of grams, the unit we use is the _____________________________________ • It is defined as one-twelfth the mass of a carbon-12 atom. o Carbon-12 chosen because of its isotope purity. • Each isotope has its own atomic mass, thus we determine the average from percent abundance. To calculate the average: • Multiply the atomic mass of each isotope by it’s abundance (expressed as a decimal), then add the results. • If not told otherwise, the mass of the isotope is expressed in atomic mass units (amu) Atomic Masses • Atomic mass is the average of all the naturally occurring isotopes of that element. Symbol Composition of the % in nature Isotope nucleus 12 Carbon – 12 C 6 protons 98.89% 6 neutrons 11
- 12. 13 Carbon – 13 C 6 protons 1.11% 14 Carbon – 14 C 6 protons <0.01% Carbon = Name _____________________________________________ Date ________________ 4.3 Section Review 1. Explain how the atomic number of an element identifies the element. 2. How can atomic number and mass number be used to find the numbers of protons, electrons, and neutrons? 3. An atom is identified as platinum – 195. a. What does the number represent? b. Symbolize this atom using superscripts and subscripts 4. How are isotopes of the same element alike? How are they different? 5. Determine the number of protons, electrons, and neutrons in each of the five isotopes of zinc. Zinc – 64: 12
- 13. Zinc – 66: Zinc – 67: Zinc – 68: Zinc – 70: 6. List the number of protons, neutrons, and electrons in each pair of isotopes. a. Li, Li b. Ca, Ca c. Se, Se 7. The atomic masses of elements are generally not whole numbers. Explain why. 8. How is the atomic mass of an element calculated from isotope data? 9. Using data for nitrogen listed in Table 5.3, (Pg 119 in your book) calculate the average atomic mass of nitrogen. Show your work. 13