The Development of Atomic Theory Chapter 4, Section 1

The Beginnings of Atomic Theory  Atoms

are EVERYWHERE!!  Determine the properties of matter.  Atomic theory was developed SLOWLY over time.  First theory of atoms was proposed over 2,000 years ago.  Greek philosopher Democritus suggested that the universe was made of invisible units…he named them atoms from the Greek word atomos which means “unable to be cut or divided”.

Democritus did not have evidence  Although

his theory explained some observations, Democritus had no evidence to support his claims.  THINK: Would you trust someone’s opinion if they had no evidence to back it up?  Even with no evidence, people supported Democritus’ claims….but other theories were also proposed.  Chemistry was developing in the 1700s and emphasis on making careful and repeated measurements in experiments was increased.  THINK: How do you think this affected Democritus’ atomic theory?

Dalton’s Atomic Theory  1808,

English schoolteacher named John Dalton proposed a revised theory  Dalton’s theory was developed on a scientific basis which (parts) hold(s) true still today.  According to Dalton, all atoms of a given element were exactly alike, and atoms of different elements could join to form compounds.

Dalton used experimental evidence  Dalton

based his theory on experimental evidence  According to the law of definite proportions, a chemical compound always contains the same elements in exactly the same proportions by weight or mass.  THINK: What does this law sound like? (something we already learned)  This helped support Dalton’s theory

Dalton’s theory did not fit all observations  Today,

Dalton’s theory is considered the foundation for modern atomic theory  As some parts turned out to be correct, some of Dalton’s thoughts could not be explained or supported with experimental evidence  Over time, the atomic theory changed as new scientists gathered more information and conducted more experiments.

Thompson’s Model of the Atom  1897,

J.J. Thompson, a British scientist, conducted an experiment that suggested that atoms were not indivisible.  Thompson was studying cathode rays when he noticed mysterious rays in vacuum tubes  His experiment suggested that cathode rays were made of negatively charged particles that came from inside atoms.  This revealed that atoms could be divided into smaller parts

Thompson developed the plumpudding model  Through

his experiments, Thompson found that when beams were deflected, this was due to charges (think magnets)  Thompson discovered electrons, which are negatively charged particles inside the atom  Thompson’s proposed model is known as the “plum-pudding model” because he thought the electrons were spread evenly throughout the atom, like blueberries in a muffin.

Rutherford’s Model of the Atom  Ernest

Rutherford, another British scientist, developed an experiment to test Thompson’s model….he found it needed to be revised.  Rutherford proposed that most of the mass of the atom was concentrated in the atom’s center

Rutherford conducted the gold-foil experiment  https://www.youtube.com/watch?v=XBqHkraf8iE

Rutherford discovered the nucleus  His

experiments suggested that an atom’s positive charge is concentrated in the center of the atom  This positively charged, dense core of the atom is called the nucleus  Data from his experiment suggested that compared with the atom, the nucleus is very small  In Rutherford’s model, negative electrons orbit that positively charged nucleus in such a way that planets orbit the sun.  THINK: What do we know about the atom today? What else is involved??

Section 1 Review  Please

complete questions 1-6 in COMPLETE SENTENCES and turn them into the basket 

The Structure of the Atom Chapter 4, Section 2

What is in an Atom?  Less

than 100 years after Dalton published his atomic theory, scientists determined that atoms consisted of smaller particles, such as the electron.  Atoms are made up of various subatomic particles  The three main subatomic particles are distinguishable by mass, charge, and location in the atom

What is in an Atom?  At

the center of each atom is a small, dense nucleus.  The nucleus is made of protons, which have a positive charge, and neutrons, which have no charge.  Moving around outside the nucleus is a cloud of very tiny, negatively charged electrons.  The mass of an electron is much smaller than that of a proton or neutron.

Each element has a unique number of protons A

hydrogen atom has one proton  A helium atoms has two protons  Each element has a unique number of protons  An element is defined by the number of protons in an atom of that element.  THINK: Why do you think protons are the only ones that matter??

Unreacted atoms have no overall charge  Even

though the protons and electrons in atoms have electric charges, most atoms do not have an overall charge.  The reason is that most atom have an equal number of protons and electrons, whose charges exactly cancel.  If an atom gains or loses electrons, it becomes charged.  A charged atom is called an ion.

The _______________ holds the atom together  THINK:

Knowing the charges of the parts of the atom, what do you think holds them together???  Positive and negative charges attract one another with a force known as the electric force.  The electric force between protons in the nucleus and electrons outside the nucleus holds the atom together

Atomic Number and Mass Number  Atoms

of different elements have their own unique structures  Because these atoms have different structures, they have different properties  Atoms of the same element can vary in structure, too  Atoms of each element have the same number of protons, but they can have different numbers of neutrons

The atomic number equals the number of protons  The

atomic number of an element, Z, tells you how many protons are in an atom of the element.  Remember most atoms are neutral because they have an equal number of protons and electrons  The atomic number also equals the number of electrons in the atom  The atomic number of a given element never changes

The mass number equals the total number of subatomic particles in the nucleus  The

mass number of an element, A, equals the number of protons plus the number of neutrons in an atom of the element  The mass number reflects the number of protons and neutrons….WHY??? What about electrons???  Although atoms of an element have the same atomic number, they can have different mass numbers because the number of neutrons can vary

Isotopes  An

isotope is an atom that has the same number of protons but a different number of neutrons relative to other atoms of the same element  Isotopes of an element have the same chemical properties  Isotopes have different masses  Isotopes of an element vary in mass because their numbers of neutrons differ  INVESTIGATE: In your book under the heading of isotopes, please find the three most common isotopes of hydrogen and how often they occur in our universe. Describe a radioisotope and what happens to them.

The number of neutrons can be calculated 

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To represent different isotopes, you can write the mass number and atomic number of the isotope before the symbol of the element If you know the atomic number and mass number, you can calculate the number of neutrons that an atom has The number of neutrons can be found by subtracting the atomic number from the mass number

Atomic Masses  The

mass of a single atom is very small  Because working with such tiny masses is difficult, atomic masses are usually expressed in unified atomic mass units.  A unified atomic mass unit (u) or (amu) is equal to one-twelfth of the mass of a carbon-12 atom 

Carbon-12, an isotope of carbon, has 6 protons and 6 neutrons, so each proton and neutron has a mass of about 1.0u…remember electrons don’t contribute to mass.

Average atomic mass is a weighted average  The

atomic mass listed for an element in the periodic table is an average atomic mass for the element as found in nature  The average atomic mass for an element is a weighted average  Commonly found isotopes have a greater effect on the average atomic mass than rarely found isotopes do

The mole is useful for counting small particles  Because

chemists often deal with large numbers of small particles, they use a large counting unit classed the mole (mol).  A mole is a collection of a very large number of particles..1 mol= 602,213,670,000,000,000,000,000 particles  Named after Amedeo Avogadro an Italian scientist…Avogadro’s number = 6.022 x 1023  THINK: Should Avogadro’s number be used for something like popcorn kernels??

Moles and grams are related  The

mass in grams of one mole of a substance is called molar mass.  In nature, elements often occur as mixtures of isotopes.  A mole of an element usually contains several isotopes  An element’s molar mass in grams per mole equals its average atomic mass in amu.

Converting between moles and grams!!   Let’s

practice!!!

Compounds also have molar mass  Recall

that compounds are made up of atoms joined together in specific ratios  To find the molar mass of a compound, you can add up the molar masses of all of the atoms in a molecule of the compound  Let’s practice!! 

Section 2 Review  Please

complete questions 1-16  On questions 1-4, 6-7, 9-12 use complete sentences  On questions 13-16 please show your work! 

Modern Atomic Theory Chapter 4, Section 3

Modern Models of the Atom  Dalton’s

theory that the atom could not be split had to be modified after the discovery that atoms are made of protons, neutrons, and electrons  Like most scientific models and theories, the model of the atom has been revised many times to explain new discoveries  In the modern atomic model, electrons can be found only in certain energy levels. Furthermore, the location of the electrons cannot be precisely predicted.

Electron location is limited to energy levels  1913,

Niels Bohr, a Danish physicist, suggested that the energy of each electron was related to the electron’s path around the nucleus  Electrons can only be in energy levels  They must gain energy to move to a higher energy level or must lose energy to move to a lower energy level  Bohr’s description of energy levels is still used by scientists today  THINK: what could be a real world example of this?

Electrons act like waves  1925,

Bohr’s model of the atom no longer explained all aspects of electron behavior.  A new model was proposed where electrons behave more like waves on a vibrating string than like particles

The exact location of an electron cannot be determined  THINK:

Looking at a spinning fan, can you precisely determine where a specific blade is located?  Determining the exact location of an electron in an atom and the speed and direction of the electron is impossible.  The best that scientists can do is to calculate the chance of finding an electron in a certain place within an atom  The darker the shading, the better the chance of finding an electron at that location…orbitals.

Electron Energy Levels  Within

an atom, electron that have various amounts of energy exists in different energy levels…and there are many levels that electrons can occupy  The number of energy levels that are filled in an atom depends on the number of electrons  The electrons in the outer energy level of an atom are called valence electrons…these electrons determine the chemical properties of an atom

Electron Transitions  The

modern model of the atom limits the location of electrons to specific energy levels  An electron is never found between these level  Electrons jump between energy levels when an atom gains or loses energy  The lowest state of energy of an electron is called the ground state  If an electron gains energy, it moves to an excited state…gains energy by absorbing a particle of light called a photon.  Electrons may lose energy and fall back to a lower energy level by releasing the photon.  THINK: What determines how much an electron will move?

Atoms absorb or emit light at certain wavelengths  The

energy of a photon is related to the wavelength of the light  High-energy photons have short wavelengths, and low-energy photons have long wavelengths  Because each element has a unique atomic structure, the wavelengths emitted depend on the particular element…giving it an “atomic fingerprint” that is unique  THINK: How could this knowledge be used in every day life? https://www.youtube.com/watch?v=kJBcXFsFa7Y

Section 3 Review  Please

complete questions 1-7 in COMPELTE SENTENCES 