Chapter 4 “Atomic Structure”

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.

Section 4.1 Defining the Atom Democritus First to suggest the existence of atoms (from the Greek word “atomos”) He believed that atoms were indivisible and indestructible. -Greek philosopher -not based on the scientific method – but just philosophy

Dalton’s Atomic Theory All elements are composed of tiny indivisible particles called atoms. John Dalton (1766 – 1844) 2) Atoms of the same element are identical. --Atoms of any one element are different from those of any other element.

Dalton’s Atomic Theory Atoms of different elements combine in whole-number ratios to form compounds. 4) In chemical reactions, atoms are combined, separated, or rearranged – but never changed into atoms of another element.

Sizing up the Atom 100,000,000 atoms = 1 cm 1,000,000 atoms = width of hair Can be observed with scanning tunneling (electron) microscopes

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.

Section 4.2 Structure of the Nuclear Atom Atoms are divisible into three subatomic particles: Electrons Protons Neutrons

Discovery of the Electron J.J. Thomson used a cathode ray tube to discover the negatively charged electron Cathode ray tubes pass electricity through a gas that is contained at a very low pressure.

Mass of the Electron Mass of the electron is 9.11 x 10-28 g The oil drop apparatus Robert Millikan determined the mass of the electron: 1/1840 the mass of a hydrogen atom

Conclusions from the Study of the Electron: Atoms have no charge, so there must be positive particles to balance the negative charge of the electrons Electrons have so little mass that other particles must account for most of the mass

Conclusions from the Study of the Electron: Eugen Goldstein observed positive proton Mass of 1 (or 1840 times that of an electron) James Chadwick confirmed the neutral neutron Mass nearly equal to a proton Protons have a positive charge. Neutrons have no charge. Mass is same.

Subatomic Particles Particle Charge Mass (g) Location Electron (e-) -1 9.11 x 10-28 Electron cloud Proton (p+) +1 1.67 x 10-24 Nucleus Neutron (no) Don’t have to know numbers, just charge and how masses compare, and location.

Thomson’s Atomic Model J. J. Thomson Thomson - plum pudding model. Electrons were like plums embedded in a positively charged pudding.

Ernest Rutherford’s Gold Foil Experiment - 1911 Alpha particles (helium nuclei) fired at a thin gold foil. Particles that hit on the detecting screen are recorded

Rutherford’s Findings Most of the particles passed right through A few particles were deflected. Conclusions: The nucleus is small, dense, and, positively charged

The Rutherford Atomic Model Based on his experimental evidence: Atom is mostly empty space. All the positive charge, and almost all the mass is in the center at the nucleus.

The Rutherford Atomic Model Nucleus is made of protons and neutrons Electrons surround the nucleus. Called the “nuclear model”

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 How then are atoms of one element different from another element?

Atomic Number Elements are different because they contain different numbers of PROTONS Atomic number - number of protons in the nucleus (smaller #) # protons = # electrons Atomic Number: # p+ : # e- : 35 35 35 53 53 53

Atomic Number 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 (C) Phosphorus (P) Gold (Au) 6 6 15 15 79 79

Mass Number Mass # = p+ + n0 Mass number is the number of protons and neutrons in the nucleus of an isotope: Mass # = p+ + n0 Atomic Number: Mass Number: # p+ : # e- : #n0 : # p+ : # e- : #n0 : 35 79.9 35 35 45 53 127 53 53 74

Mass Number Practice Atom p+ n0 e- Mass # 16 8 8 8 33 41 33 74 15 16 Oxygen 16 8 8 8 Arsenic 33 41 33 74 Phosphorus 15 16 15 31

Complete Symbols Contain the symbol of the element, the mass number and the atomic number. Atomic number X Superscript → Subscript → Mass number *

Na Symbols 11 11 12 23 11 11 23 Find each of these: number of protons number of neutrons number of electrons Atomic number Mass Number 11 11 12 Na 23 11 11 23 *

Symbols If an element has an atomic number of 34 and a mass number of 78, what is the: number of protons = number of neutrons = number of electrons = complete symbol 34 43 34 34 X 78 *

If an element has 91 protons and 140 neutrons what is the Symbols If an element has 91 protons and 140 neutrons what is the Atomic number = Mass number = number of electrons = complete symbol 91 131 91 *

If an element has 78 electrons and 117 neutrons what is the Symbols If an element has 78 electrons and 117 neutrons what is the Atomic number Mass number number of protons complete symbol *

Isotopes Dalton was wrong about all elements of the same type being identical Atoms of the same element can have different numbers of neutrons. Thus, different mass numbers. These are called isotopes. Atomic #: Mass #: # p+: #n0: Atomic #: Mass #: # p+: #n0: Atomic #: Mass #: # p+: #n0: *

Isotopes Isotopes are atoms of the same element with different masses, due to varying numbers of neutrons.

We can also put the mass number after the name of the element: Naming Isotopes We can also put the mass number after the name of the element: carbon-12 Mass: carbon-14 Mass: uranium-235 Mass: 12 14 235 *

What’s the only thing that changes? # of neutrons Isotopes are atoms of the same element having different masses, due to varying numbers of neutrons. Isotope Protons Electrons Neutrons Nucleus Hydrogen–1 (protium) 1 Hydrogen-2 (deuterium) Hydrogen-3 (tritium) 2 What’s the only thing that changes? # of neutrons

Atomic Mass How heavy is an atom of oxygen? Depends - there are different masses of oxygen atoms. We want the average atomic mass. Based on abundance (percentage) of each variety of that element in nature. *

Measuring Atomic Mass Measure atomic mass with the Atomic Mass Unit (amu) Defined as one-twelfth the mass of a carbon-12 atom. Each isotope has its own atomic mass, thus we determine the average from percent abundance. We don’t use grams for this mass because the numbers would be too small. Carbon-12 chosen because of its isotope purity. – 6 p+, 6 n0 almost all mass and high % C12. *

To calculate the average: Multiply the atomic mass of each isotope by it’s abundance (expressed as a decimal), then add the results. Expressed as amu. C-12 = 12 amu. *

Composition of the nucleus Atomic Masses Atomic mass is the average of all the naturally occurring isotopes of that element. Isotope Symbol Composition of the nucleus % in nature Carbon-12 12C 6 protons 6 neutrons 98.89% Carbon-13 13C 7 neutrons 1.11% Carbon-14 14C 8 neutrons <0.01% Carbon = 12.011 *

Atomic Mass Example (11.0)(.802) 10.8 amu (10.0)(.198) B-10 = 19.8% At. Mass = + = (10.0)(.198) (11.0)(.802) 10.8 amu

The Periodic Table: A Preview Periodic table - arrangement of elements in which the elements are separated into groups based on a set of repeating properties. Allows easy comparison of the properties of different elements

The Periodic Table: A Preview Period - horizontal row (there are 7 of them) Group - vertical column Also called a family Elements in a group have similar chemical and physical properties Identified with number and “A” or “B”

Draw an arrow and label a period and a group.