--- ## Sections - All elements want to be in their most *stable state* which is the represented by the **noble gas** in their respective row. - Alkali metals are *extremely reactive* increasing in reactivity going down the periodic table. - Halogens and other **non-metals** are heavily used in organic chemistry since they are fairly stable and abundant. - See [Atomic structure] for information on how naming and each element is formatted. > [! Periodic Table] > > ![[Pasted image 20230110123606.png | 850]] > - Groups are columns and periods are rows. > > - Depending on how far away a period is from a noble gas the greater the charge of elements in that row. *Right of the table is positive charge left is negative*. --- ## Trends - **Ionization energy** is the **energy** required to remove one of the outer most [[Electrons]] from an element. Or in other words the capability of an element to enter into a reaction that requires it to donate [[Electrons]] forming an [[Ion]]. >[!tip] >#### More on ionization energy >- The ionization energy will be lower the **less valence electrons an atom has**. >- The second ionization energy is the energy required to take the **second electron off of an atom**. >- If the 3rd ionization energy is higher than the 1st and 2nd that means that **the third atom is in another fully filled shell** which assumes the role of the [[Valence Electrons|valence shell]] after the 1st and 2nd electrons were removed *(in a molecule that has 2 valence electrons). * - **Electron affinity** is the amount of energy gained when an [[Electrons]] binds to that atom to from a *negatively* charged [[Ion]]. The greater the electron affinity the more easy it is for an element to *accept* electrons. - **Electronegativity** is very similar to this but instead describes how **much an atom wants another electron.** Not the amount of energy released. Ionic bonds have large differences in electronegativity because one element want an electron while the other does not. *Makes sense right?* - **Metallic character** is loosely defined as the **reactivity** of an element. This is a *result* of the increasing ionization energy and electron affinity of elements. - **Atomic radius** is the size of the atom **including** the electrons that surround it. >[!tip] >- The size of an atom increases towards the bottom left. >- Positive ions *cations* are smaller than neutral atoms because shielding decreases while nuclear charge stays the same. Anions are larger *under same logic.* >[!caution] >#### Math moment with these numbers: >![[Periodic table example 3.png |750]] > [! Atomic Trends] > > <br> > >![[trends.png]] >- The number of [[Valence Electrons]] in an element is denoted by the column *(from left to right not including the shorter metallic columns)* number it is in. > >- Elements are grouped in **columns** of other elements that share physical properties. --- ## Metals and their properties - Metals are efficient conductors of **heat** and **electricity**. - Metals are very malleable *(can be hammered into sheets)* and ductile *(They can be pulled into wires)*. - Most metallic atoms tend to *lose* electrons to form **Positive [[Ion |ions]]**. - Metals are good conductors of heat so **if one side of them is heated the other side will also be hot**. --- ## Non-metals and their properties - Non metals tend to gain electrons to form **anions** in reactions with **metals** since they have too few electrons in their [[Valence Electrons|valence shell]]. Remember however, that this is only true for non-metals on the right of the periodic table. - Often bond to each other to form **[[Covalent bond |covalent bonds]].** --- ## Blocks - The periodic table has what are referred to as different *blocks* ranging from s to f which come from the orbital subshell letters. > [! Arrangement of blocks in periodic table] > <br> > > ![[Periodic table blocks.png]] --- #mainpage