Section 2.1 Atoms and Their Structures

Relate historical experiments to the development of the atom, Illustrate the modern model of an atom, Interpret the information available in an element block of the periodic table

Solving a problem: 1) Observation - use senses to observe the behavior of matter 2)HYPOTHESIS – testable prediction Hypotheses that are verified by repeating experiments

EXPERIMENT: Investigation (with a control) designed to test a hypothesis Hypotheses lead to scientific theories THEORY: Explanation based on many observations and supported by the results of many experiments.  Ex: Dalton’s Atomic Theory

SCIENTIFIC LAW: Fact of nature that is observed so often that it becomes accepted as truth. A law can be used to make predictions, but does not explain why something happens Example: Sun rises in the east Theories can explain laws

Greek Philosophers (2500 years ago)  4 Fundamental Elements- air, earth, fire, and water Questioned if matter could be divided endlessly into smaller pieces or if there was an ultimate small particle of matter Aluminum Foil

Proposed the world is made up of empty space and the smallest particles of matter are called ATOMS  This introduced the atomic theory of matter  Different types of atoms exist for every type of matter

Antoine Lavoisier (Luh-voh-zee-ay) 1782  Concluded that when a reaction occurs, matter is neither created nor destroyed  Law of Conservation of Matter

Atoms are neither created nor destroyed You can’t throw anything away because there is no “away” Recycling Nitrogen- Nitrogen is atmosphere is converted into compounds used on Earth, then returned to atmosphere (p. 53, Figure 2.4)  Recycling plastic, aluminum and glass- reusing atoms in these materials, we imitate nature and conserve natural resources (in natural processes atoms are recycled)

Observed that the elements that composed compounds were always in a certain proportion by mass- LAW OF DEFINITE PROPORTIONS Ex: water is always 11% H and 89% O by mass

Theory essentially intact with small modifications to accommodate new discoveries Main Points of Dalton’s Atomic Theory 1. All matter is made up of atoms 2. Atoms are indestructible and cannot be divided into smaller particles. 3. All atoms of one element are exactly alike, but atoms are different for different elements Late 19 th century, experiments began to suggest that atoms are made up of even smaller particles (electrons, protons, neutrons)

Today we know that atoms are made of smaller particles and Atoms of the same element can be nearly the same (but not exactly) Cathode Ray Tube Experiment- J.J. Thomson (1897) Vacuum tube- positive and negative electrode- ray travels through tube from the – to the +, rays bent toward + and away from the – Electrons, Protons, Neutrons ELECTRON- negatively charged subatomic particle Mass equal to 1/1837 the mass of a Hydrogen atom PROTON- positively charged subatomic particle The amount of charge on an electron and a proton is equal and opposite The mass of a proton is much greater than the mass of an electron (slightly less then a Hydrogen atom)

Until 1910, it seemed that atoms were made up of equal numbers of electrons and protons J.J. Thomson discovered that Neon consisted of atoms of two different masses

Ernest Rutherford “Gold Foil Experiment” (p.62, Fig 2.9) + charged particles (alpha particles) sent through a gold foil – most went straight through (empty space), few deflected (hit nucleus) Revealed the arrangement of the atom Atom is nearly all empty space with a small, dense, positively charged core called a NUCLEUS erford/

Atoms of an element that are chemically alike but different in mass are called ISOTOPES The discovery of isotopes – atoms must contain a third type of particle that explains mass differences NEUTRON- neutral subatomic particle Mass is equal to that of a proton but has no electrical charge Existence of the neutron was confirmed in early 1930s

Quarks – small particles of matter that make up protons and neutrons.  6 “flavors” or types – top, bottom, charm, strange, up and down  An arrangement of 3 of these will form a proton, another arrangement will form a neutron.

 Subatomic particles produced by the decay of radioactive elements  Elementary particles that lack an electric charge  F. Reines would say, "...the most tiny quantity of reality ever imagined by a human being

 Copiously produced in high-energy collisions  Traveling essentially at the speed of light  Unaffected by magnetic fields neutrinos Their unique advantage arises from a fundamental property: they are affected only by the weakest of nature's forces (but for gravity) and are therefore essentially unabsorbed as they travel cosmological distances

ATOMIC NUMBER: Number of protons in the nucleus of an atom of that element # of protons determines the identity of an element and many of its chemical and physical properties Atomic number also tells us the number of electrons in a neutral atom of an element (p + = e - )

MASS NUMBER: Sum of the protons and neutrons in the nucleus of an atom Mass # = p + + n 0 n 0 = Mass # - p + Example: # of neutrons: mass # = 19 (Flourine) n 0 = = 10

Isotopes have different mass numbers because they have different numbers of neutrons (they have the same atomic number) Isotopes are identified by placing the mass number after the name or symbol of the element Ex: Li-7, Li-6 Ne-20, Ne-21, Ne-22

 Each box contains: element, state, atomic number, symbol and average atomic mass (weighted average of the naturally occurring isotopes) Atomic mass example: Chlorine has 2 isotopes- Cl-37 and Cl % is Cl-37 and 75.8% is Cl-35 So the atomic mass is 35.45

Protons, Neutrons, Electrons Particle Symbol Charge Mass in grams Mass in u Proton p x Neutron n o x Electron e x u – ATOMIC MASS UNIT (devised mass unit)  1 u = 1/12 the mass of a carbon-12 atom  1 u  mass of single proton or neutron