Atomic Structure Chapter 4

Section Overview 4.1: Defining the Atom 4.2: Structure of the Nuclear Atom 4.3: Distinguishing Among Atoms

Defining the atom Section 4.1

Early Models of the Atom An atom is the smallest particle of an element that retains its identity in a chemical reaction. Democritus’s Atomic Philosophy: Believed that atoms were indivisible and indestructible. However, his idea lacked explanation of chemical behavior and experimental support Dalton’s Atomic Theory: Transformed Democritus’s ideas on atoms into a scientific theory using experiments relying on the scientific method. All elements are compose of atoms. Atoms of the same element are identical. Atoms of different elements are different. Atoms of different elements can mix together or chemically combine to form compounds. Chemical reactions occur when atoms are separated, joined, or rearranged.

Sizing up the Atom Despite their small size, individual atoms are observable with instruments such as scanning tunneling microscopes. Individual atoms can even be moved around and arranged into patterns. This holds promise for thee creation of atomic-sized electronic deices. This is called the “nanoscale” and could become essential to future applications in medicine, communications, solar energy, and space exploration.

Structure of the nuclear atom Section 4.2

Subatomic Particles One important change to Dalton’s theory today is that atoms are now known to be divisible. They can be broken down into three types of subatomic particles: electrons, protons, and neutrons. Electrons: Negatively charged subatomic particle (discovered by J. J. Thomson). The mass (discovered by Robert A. Millikan) is 1/1840 the mass of hydrogen. Protons: Positively charged subatomic particles (discovered by Eugen Goldstein) with a mass about 1840 times that of an electron. Neutrons: No charge (discovered by James Chadwick) with a mass almost equal to a proton. *The net charge of an atom is neutral*

The Atomic Nucleus After the discovery of the subatomic particles, many scientists wondered how these were arranged inside of an individual atom. Rutherford’s Gold-Foil Experiment: Set up an experiment where he directed a narrow beam of alpha particles at a thin sheet of gold foil. The hypothesis was that the alpha particles would go straight through with only a slight deflection. However, the great majority passed through without any deflection with some bouncing straight back. Rutherford’s Atomic Model: His experiment led to a new model of the atom in which the atom is mostly empty space (explaining lack of deflection) with a nucleus in the center. * In the atomic atom, protons and neutrons are located in the nucleus with the electrons distributed around it occupying most of the volume*

Distinguishing among atoms Section 4.3

Atomic Number Elements are different because they contain a different number of protons. The atomic number of an element is the number of protons in the nucleus of that atom (ex. All hydrogen atoms have one proton, so their atomic number is one). Since the net charge of an atom is neutral, the number of protons are equal to the number of electrons.

Mass Number The mass of an atom is concentrated in the nucleus and therefore depends on the number of protons and neutrons. The total number of protons and neutrons an atom is called the mass number (ex. Helium has two protons and two neutrons and therefore a mass number of 4). The number of neutrons in an atom is the difference between the mass number and the atomic number. The composition of an atom is written using the symbol, the atomic number, and the mass number (ex. 197Au). 79

Determining the Composition of an Atom Example Problem: How many protons, electrons, and neutrons are in each atom? Beryllium (Be) atomic #=4 mass #=9 Neon (Ne) atomic #=10 mass #=20 Sodium NA) atomic#=11 mass#=23

Determining the Composition of an Atom Example Problem: How many protons, electrons, and neutrons are in each atom? Beryllium (Be) atomic #=4 mass #=9 Solution: Knowns Unknowns Atomic # # of protons # of neutrons Mass # # of electrons Remember: # of protons = atomic # and # electrons=# of protons So: 4 protons and 4 electrons Remember: # of neutrons = mass # - atomic # So: 9-5 = 5 neutrons

Determining the Composition of an Atom Example Problem: How many protons, electrons, and neutrons are in each atom? Neon (Ne) atomic #=10 mass #=20 Solution: Knowns Unknowns Atomic # # of protons # of neutrons Mass # # of electrons Remember: # of protons = atomic # and # electrons=# of protons So: 10 protons and 10 electrons Remember: # of neutrons = mass # - atomic # So: 20-10 = 10 neutrons

Determining the Composition of an Atom Example Problem: How many protons, electrons, and neutrons are in each atom? Sodium NA) atomic#=11 mass#=23 Solution: Knowns Unknowns Atomic # # of protons # of neutrons Mass # # of electrons Remember: # of protons = atomic # and # electrons=# of protons So: 11 protons and 11 electrons Remember: # of neutrons = mass # - atomic # So: 23-11 = 12 neutrons

Isotopes Sometimes, atoms of the same element do differ. Isotopes are atoms that have the same number of protons but different numbers of neutrons. Because isotopes have a different number of neutrons, they all have a different mass number. Isotopes are still chemically alike because they still have the same number of protons and electrons.

Atomic Mass Scientists use a mass spectrometer to determine the mass of atoms. Although this can be useful, it is more useful to compare the relative masses of atoms using a reference isotope as a standard. The standard used is carbon-12. An atomic mass unit (amu) is defined as one twelfth of the mass of a carbon-12 atom. The atomic mass of an element is a weighted average mass of the atoms in a naturally occurring sample of that element which takes into account isotopes. To calculate the atomic mass, multiple the mass of each isotope by its natural abundance, expressed as a decimal, and then add the products.

Atomic Mass Example Problem: Element X has two natural isotopes. The isotope with a mass of 10.012 amu has a relative abundance of 19.91%. The isotope with a mass of 11.009 amu has a relative abundance of 80.09%. Calculate the atomic mass of this element.

Atomic Mass Example Problem: Element X has two natural isotopes. The isotope with a mass of 10.012 amu has a relative abundance of 19.91%. The isotope with a mass of 11.009 amu has a relative abundance of 80.09%. Calculate the atomic mass of this element. Solution: Knowns Unknown Isotope X-10 Isotope X-11 atomic mass of X Mass= 10.012 amu Mass= 11.009 amu Abundance=80.09% Abundance=19.91% 10.012 amu x 0.1191 = 1.993 amu 11.009 amu x 0.8009 = 8.817 amu 1.993 amu + 8.817 amu = 10.810 amu

The Periodic Table - Preview A periodic table is an arrangement of elements in which the elements are separated into groups based on a set of repeating properties. A periodic table allows you to easily compare the properties of one element (or a group of elements) to another element (or group of elements). The atomic number is above each element symbol and the atomic mas is below each symbol. Each horizontal row is called a period. Each vertical column is called a group. Elements in the same group have similar chemical and physical properties.

The Periodic Table - Preview