periodic Table Flashcards
learn these (36 cards)
matter
Anything that has mass, all objects have matter.
atoms.
The basic building blocks of all matter, can combine to form molecules but cannot be divided into smaller parts by ordinary chemical means.
Have the same number of protons as electrons so the whole atom has no electric charge.
element
A substance made up of only 1 type of atom, each with the same number of protons.
structure of the atom- nucleus, electrons, protons, neutrons.
Nucleus is surrounded by orbiting electrons which are negatively charged and arranged in shells.
The nucleus, the center of the cell, contains protons which are positively charged and neutrons which are negatively charged.
Protons have the same mass as neutrons.
Atomic particle, relative mass and relative charge.
Proton, 1, +1
Neutron,1,0
electron,1/1840,-1
History of the atom, model dalton
Dalton’s Model (1803)
Proposed the atom as a solid, indivisible sphere.
Atoms of different elements vary in mass and properties.
Introduced the concept of chemical compounds formed from atoms.
History of the atom, model Thomsons
Thomson’s Model (1897)
Discovered the electron through cathode ray experiments.
Proposed the “Plum Pudding Model,” suggesting atoms are spheres of positive charge with embedded electrons.
Introduced the idea of subatomic particles.
History of the atom, model Rutherford.
Rutherford’s Model (1911)
Conducted the gold foil experiment, revealing a dense, positively charged nucleus.
Proposed that electrons orbit the nucleus, leading to the “Nuclear Model” of the atom.
Introduced the concept of empty space within the atom.
History of the atom, Bohr.
Bohr’s Model (1913)
Built on Rutherford’s model by introducing quantised energy levels for electrons.
Suggested electrons move in fixed orbits around the nucleus with specific energy levels.
Explained atomic spectra and stability of electron orbits.
Benefits of models in science.
Makes complex things easier to understand.
Provides clear visual representations.
Helps predict outcomes and guide experiments.
Aids in experimental design and variable identification.
Useful across various scientific fields.
Limitations of models in science.
Can lead to inaccuracies.
Rely on assumptions that may not always hold true.
May become outdated with new data.
Choice of model can introduce bias.
Often miss key variables and interactions.
Atoms of different elements differ in num. of protons (atomic number)
The atomic number of an element tells the number of protons in the nelcused of that element.
e.g. lithium atomic number = 3, nitrogen atomic number =7
Use periodic table to determine atomic number
The atomic number, or number of protons is the bottom of the two numbers next to the symbol.
Use periodic table to determine mass number
The mass number which is the number of protons and neutrons is the top number next the symbol.
Use periodic table to determine the element symbol.
The symbol of the element is the letters on the periodic table.
Use mass and atomic number to determine number of neutrons.
Number of protons is the atomic number.
So its the mass number minus the atomic number for the amount of neutrons.
draw diagrams to represent the nuclear and electronic configuration for the
first 20 elements., how to make them.
Electrons are arranged in shells around the nucleus.
Electrons fill the first shell, then second then third and so on.
First shell can fit 2, the second 8, the third 18 and the fourth 32.
Calculated using the formula 2n squared, n the number of the electronic shells.
Number of electrons are the same as the number of protons, which is the atomic number so the atomic number will also tell the amount of electrons.
Position of metals and non-metals on periodic table
First 100 elements are arranged in order of increasing atomic number.
Then arranged in row so elements with similar properties are in the same column.
Metals can be found mainly in Groups 1 and 2 in the central block. And non metals in groups 7 and 8.
identify the names of some of the
chemical groups on the Periodic Table
Group 7: Halogens
Characteristics:
Have 7 electrons in their outer shell.
Reactivity decreases down the group as the outer electrons get further from the nucleus.
Examples: Fluorine (F), Chlorine (Cl), Bromine (Br).
Group 8: Noble Gases
Characteristics:
Have full outermost shells (8 electrons, except Helium which has 2).
Extremely stable and do not need to gain or lose electrons, making them largely unreactive.
Examples: Helium (He), Neon (Ne), Argon (Ar).
Alkali Metals (Group 1)
Characteristics: Highly reactive, especially with water; soft metals; have one valence electron.
Examples: Lithium (Li), Sodium (Na), Potassium (K).
What is ionic bonding.
- where electrons completely transfer.
- occurs between metals and non-metals
- involves a transfer of electrons form one atom to another to form electrically charged “atoms” called ions.
- each with a complete outer electron shell, so very unreactive like nobel gasses.
Example of Ionic Bonding
e.g. Sodium and Chlorine to form sodium chloride.
- sodium has one electron in its outer shell
- transfers to chlorine atom, both have now 8 electrons in the outer shell.
- now the atoms are ions and NaCI.
Properties of Ionic Bonding.
- high melting points
- conduct electricity when molten or in a soluction bc the charged ioins and free to move around.
sometimes dissolve in water
crystalline.
What is covalent bonding.
- When two nonmetals combine, the atoms share electrons resulting in the formation of a covalent bond and a covalent compound.
- The idea of stability is the same, resulting when each atom has eight
elecgrons in its outermost orbit.
-results in formation of molecules
Examples of covalent bonding.
e.g. fluorine reacting with fluorine to form fluorine gas.
- When they meet they share one electron with each toerh, so there are 8 in its outermost orbit, stable molecule (diatomic molecule) is formed.
- contains 2 fluorine atoms joined together by a single covalent bond(formed when one pair of electrons is shared)
-valency is still 1.
